
Class QT . 

Book Jll4S 

Copight]\ 

CQPXRIGHT DEPOSm 



ELEMENTARY CHEMISTRY 



FOR HIGH SCHOOLS AND ACADEMIES 



ALBERT L. AREY 

ROCHESTER (N.Y.) HIGH SCHOOL 



ftgT^ This pamphlet is sold at the retail price of 
Arey's " Chemistry," and can be exchanged at 
any bookseller's about November 1 9 without fur- 
ther charge, for a copy of the complete work. 



T$zto fgork 
THE MACMILLAN COMPANY 

LONDON : MACMILLAN & CO., Ltd. 

1899 

All rights reserved 




ELEMENTARY CHEMISTRY 



■J*&V$o- 



ELEMENTARY CHEMISTRY 



FOR HIGH SCHOOLS AND ACADEMIES 



" f 
ALBERT L. AREY 

ROCHESTER (N.Y.) HIGH SCHOOL 



THE MACMILLAN COMPANY 

LONDON : MACMILLAN & CO , Ltd. 
1899 

All rights reserved 



61395 

COPTBieHT, 1-'.)'.' 



By THE MACMILLAN COMPANY 



SECOND COPY, 



NoriuooD ^reas 

J. S. Cuahing & Co. - Berwick & Sir 



Norwood Mass. U.S.A. 






CHAPTER I 
CHEMICAL ACTION 

1. Elements and Compounds. — The different kinds of 
natter known to man may be divided into two classes : 

(1) Compounds, or those which may be decomposed or 
eparated into other substances. 

(2) Elements, or those which have thus far resisted all 
attempts to decompose them. 

About seventy simple substances or elements have been 
liscovered, and, so far as is at present known, all com- 
pounds are the result of the chemical union of two or more 
3f these. 

It is possible that, as our knowledge of chemistry in- 
3reases, many, if not all, of the substances now classed as 
elements may be shown to be compounds. Water was con- 
sidered an element until 1783, and several other so-called ele- 
ments have been resolved into simpler forms since that time. 

There are many chemists who consider the seventy ele- 
ments as so many unsolved problems. 

2. Molecules and Atoms. — The accepted theory of the 
constitution of matter maintains : — 

1. That it is made up of minute particles called mole- 
cules (little masses), each one of which, in a given sub- 
stance, is exactly like its neighbors in weight, volume, and 
structure. 

2. That they move about each other, under the influence 
of heat, as separate bodies. 



2 CHEMISTRY 

3. That they are the limit beyond which it is impossible 
to subdivide matter without destroying its identity. 

In accordance with this theory, each molecule of a com- 
pound is believed to contain the same elements that chemi- 
cal analysis shows the large masses of the substance to 
contain, and these smaller portions of the elements are 
called atoms. There is good reason for believing that 
atoms rarely exist in a free state, but that the molecules 
of most elements consist of two or more atoms. 

A molecule is the smallest particle of a substance which can 
exist in the free state, and which has the same composition as 
din/ larger mass of the same substance. 

.1// atom is tin* smallest particle of an element that exists in 
(in'/ molecule. 

We may now state the following definitions : — 

.1 compound is a substance whose molecule contains two or 
more kinds of atoms. 

.1// element is a substance ichose molecule contains only one 
kind of atoms. 

3. The Domain of Chemistry. — Chemistry is that branch 
of science which deals with, changes in the identity of sub- 
stances; and ivith the laws, causes, and effects of such changes. 

The subject is closely related to physics ; every chemical 
change is accompanied by some physical change, but the 
chemical change differs in one important particular from a 
physical change: the chemical change is due to forces act- 
ing upon atoms, while the physical change depends upon 
forces acting upon the molecule. 

A physical change is one which does not destroy the identity 
of the substance acted upon. 



CHEMICAL ACTION 3 

Illustration. When a bar of steel is magnetized it acquires 
a new property, but it remains the same substance, and the 
change is physical. 

A chemical change is one which destroys the identity of the 
substance acted upon. 

Illustration. When a bar of steel rusts a portion of the 
steel is converted into a new substance which differs from 
the steel in color, tenacity, elasticity, and other properties. 

4. Chemical Action. — In some instances one may be in 
doubt as to whether a chemical change has taken place, and 
in a few instances chemical analysis is necessary to prove it. 
But in general the occurrence of any of the following phe- 
nomena, when two or more substances are mixed, may be 
taken as evidence of chemical action : — 

1. Effervescence. 

2. The evolution of heat and light. 

3. Change of color. 

4. Change of volume. 

5. Change of state. 

6. The development of electricity. 

Exceptions. 1. A change of state by solution of a solid 
or gas. 

2. A change of volume due to the absorption of a gas 
by a solid or liquid, or to- a change in temperature. 

Take notes on the following experiments which will be 
performed for you, and designate them as physical or 
chemical changes : — 

Experiment I. — Sugar and potassium chlorate are mixed, and a 
drop of sulfuric acid added. 

Experiment II. — Sulfuric acid is added to syrup. 

Experiment III. — A rubber ruler is electrified. 

Experiment IV. — A beam of sunlight is decomposed with a prism. 



CHEMISTRY 



Experiment V. — A piece of platinum wire is heated to redness. 

Experiment VI. — A piece of magnesium wire is heated to redness. 

Experiment VII. — Sulfur and potassium chlorate are mixed in a 
mortar with considerable friction. 

Experiment VIII. — Solutions of potassium iodid and mercuric 
chlorid are mixed. 

Experiment 9 and the succeeding experiments are to 
be performed by the pupil unless special directions to the 
contrary are given. 

Experiment IX. A Chemical Change. — 1. Examine a piece of 
marble carefully, fix its appearance in mind, so that you can detect 
any change. 

2. Drop a small piece in the test bottle and cover it with dilute 
hydrochloric acid. What occurs ? 

3. After a short time test the gas in the upper part of the test 
bottle with a lighted match. Does the match continue to burn ? Is 
the gas combustible? Is the gas ordinary air? Why? Does the 
marble disappear '.' 

4. In order to tell whether the marble has been changed chemically, 
the acid must be expelled. To accomplish this, pour the solution into 
an evaporating dish, place it on a piece of wire gauze, and bring the 
liquid to a boil. When the liquid begins to solidify and turn yellow, 
add a few drops of water, repeating, if necessary, to obtain a solid 
white residue. 

Examine this residue, compare it carefully with marble. Set the 
residue aside for 24 hours to determine whether it is permanent when 
exposed to the air. 

Fill out the following table : — Marble Residue 



Is it hard or soft ? 

Does it effervesce with hydrochloric acid ? 

Is it sol ul »le in water ? 

Is it permanent in air ? 

How do the properties of the residue compare with those of marble ? 

Is it marble ? 

What have you proven ? 

Experiment X. — Bring together on a flower-pot saucer a little 
phosphorus and iodin. What evidence have you that chemical 
action took place ? Have either of the original substances disap- 
peared ? Has a new substance been formed ? It will be seen that 
simple contact is sufficient to cause the two substances to act upon 
each other. 



CHEMICAL ACTION 5 

Does either substance melt ? Why ? 

Is this a case of chemical action between solids ? 

Is the action as energetic at first as it is after a few seconds ? 
Explain. 

Cautiox. — Handle phosphorus with great care ; it takes fire when 
rubbed or cut in the air, and should always be kept in water. 

5. Conservation of Matter. 

Experiment XL — Pour 10 cc. of dilute sulfuric acid into a beaker. 
In a second beaker pour an equal quantity of calcium chlorid solution. 
Place both beakers in one scale pan and balance them carefully with 
weights, sand, or shot, placed in the other scale pan. Now pour the 
calcium chlorid into the sulfuric acid. Does a chemical change occur ? 
Replace the beakers and determine whether the weight of the beakers 
and their contents has been changed. Does chemical action change 
the total quantity of matter in existence ? Was the total quantity of 
sulfuric acid in the world increased or diminished by the above ex- 
periment ? How was the total quantity of calcium chlorid affected ? 
of the white substance formed ? 

6. The Effect of Solution on Chemical Action. 
Experiment XII. — Place as much baking soda as you can take on 

the end of a knife blade in a dry test bottle. Add an equal amount 
of tartaric acid ; shake the bottle to mix the powders thoroughly. 
Has any change occurred ? 

Pour a few cubic centimetres of water into the bottle. What evi- 
dence of chemical action do you observe ? 

Does solution aid chemical action ? Is it because more intimate 
contact of the molecules is obtained when solutions are mixed than is 
possible with solids ? Should diminishing cohesion assist chemical 
action ? State your opinion as to why solution aids chemical action. 

7. Effect of Heat on Chemical Action. 

Experiment XIII. — 1. Mix six grammes of potassium chlorate and 
one gramme of powdered charcoal thoroughly. What occurs ? 

2. Apply a lighted match. Was the change chemical or physical ? 

How does the operation of striking a match illustrate the effect of 
heat upon chemical action ? 

Why do metals rust more rapidly when hot than at lower tempera- 
tures ? Experiment 22 illustrates this effect. Do fuels combine with 
the air when cold ? 



6 CHEMISTRY 

8. Light causes Chemical Action. 

Experiment XIV. — Cat a design from tin-foil and place it on a 
piece of blue print paper. Expose paper and design to sunlight for 
a few minutes. Wash the paper in water. 

Has the sunlight affected the exposed chemical ? In what way ? 

The art of photography is based on the action of light on 
chemicals. In growing plants sunlight causes the decompo- 
sition of carbon dioxid, which is only accomplished by the 
chemist with difficulty. 

In the preparation of hydrochloric acid by synthesis de- 
scribed on page 92, the chemical action is assisted by light. 

Query. — Why do certain colors fade when exposed to light ? 

9. Pressure. — When the two gases, hydrochloric acid 
and hydrogen phosphid, are subjected to increasing press- 
ure they combine to form a crystalline solid known as 
phosphonium chlorid. Similarly sulfur and powdered lead 
may be caused to combine by great pressure, forming lead 
sulnd. 

Quebt. — What relation does this action suggest between the inten- 
sities of chemical affinity and distances between molecules ? 

10. Concussion or Detonation. — In a very few cases, 
chemical action is brought about by detonation. The 
molecules of the gas acetylene consist of two atoms of 
carbon united with two of hydrogen. If a small quantity 
of mercury fulminate be detonated near a globe filled with 
this gas the carbon is instantly deposited in solid form and 
the hydrogen liberated. This action is not fully under- 
stood; some chemists believe that the particular form of 
sound vibration produced disturbs the motions of the atoms 
constituting the molecule, and thus causes disruption. 

11. Electricity. — If a current of electricity be jessed 
through a solution of copper sulfate, the compound is de- 



CHEMICAL ACTION 7 

composed, and many other compounds are affected in the 
same way. In Experiment 40 this effect is also illustrated. 

12. The Effect of Trituration on Chemical Action. 
Experiment XV. — Using pincers, hold a small lump of rosin in the 

Bunsen burner flame, observe the character of the flame produced by 
the burning rosin. Is rosin easily ignited ? Does it burn rapidly ? 
Does the rosin melt before it ignites ? 

Experiment XVI. — Triturate a small piece of rosin in a mortar, 
fill the end of a large glass tube with the powder and blow it into the 
burner flame. Does the finely divided rosin burn with a smoky flame 
or does it flash ? Does it burn as rapidly as in the previous experi- 
ment ? How does the energy of the chemical action compare with 
that observed in the last experiment ? In which case is the higher 
temperature reached ? 

Experiment XVII. — Make a compact pile of about \ cu. cm. of 
powdered rosin on a piece of porcelain or earthenware, ignite with a 
Bunsen burner. How does the chemical action compare with that of 
the previous experiment ? Does the increased chemical action depend 
upon the size of the particles ? Would a solid piece having the same 
area as the sum of the surfaces flash ? Does the chemical activity 
depend upon the surface only ? Upon the mass of the particles only ? 

13. Mechanical Mixture. 

Experiment XVIII. — 1. Mix about four grammes of sulfur and an 
equal weight of fine wrought iron filings on a sheet of paper. Divide 
into three portions. 

2. Examine the first portion with a magnifying glass. Can you 
distinguish the particles of sulfur from those of iron ? Can you 
separate the iron from the sulfur with a magnet ? Now put the mix- 
ture in a test tube and pour water on it. Are the substances combined 
or not ? Shake the tube ; what is the yellow substance floating on 
the water ? Has chemical action taken place ? 

3. Treat the second portion with carbon disulfid. What is the 
black substance at the bottom of the tube ? What has happened ? 
Is the color of the carbon disulfid changed ? What does this indicate ? 
Is a chemical compound formed in this experiment ? 

Experiment XIX. — 1. Put the third portion of the mixture made 
in Experiment 18, in a dry test tube and heat gently. When it is red 
hot remove the tube from the flame. Is there any evidence of com- 
bustion in the tube ? 



8 CHEMISTRY 

2. After the action is over and the tube has cooled down, loosen 
the contents with a short piece of wire, and pour it out on a piece of 
paper. Does the mass look like the mixture of sulfur and iron with 
which you started ? 

3. Examine with a magnifying glass. Can you separate the sulfur 
and iron with water as before ? Can you separate them with a 
magnet ? 

4. Treat a portion of the mass with carbon disulfid. Is the effect 
the same as before ? Is the color of the carbon disulfid changed ? 
What do you conclude concerning the effect of heat on the mixture ? 

REVIEW QUESTIONS 

1. Define chemical action. What assists it ? What retards it ? 

2. Mention those conditions which aid chemical action, (a) by 
decreasing the distance between the unlike molecules, (6) by dimin- 
ishing the cohesion of the factors. 

3. Describe an experiment to show that there is no loss of matter 
in chemical change. 

4. Distinguish between a mechanical mixture and a chemical com- 
pound. Illustrate each. 

5. Distinguish between chemistry and physics ; between atoms and 
molecules ; between chemical changes and physical changes. 

6. What mechanical mixture was formed in Experiment 12 ? 
In what part of the experiment were chemical compounds formed ? 

Write answers to these questions in your note-book. 



CHAPTER II 
SYMBOLS AND LAWS 

14. Symbols. — Chemists of all countries have agreed to 
use the initial letter of the Latin name of an element as an 
abbreviation which shall stand for a single atom of that 
element. In case two or more elements begin with the 
same letter the second characteristic letter is added to the 
symbol, thus : — 

C Carbon N Nitrogen Na* Sodium 

Ca Calcium S Sulfur K* Potassium 

CI Chlorin Si Silicon Ag Silver 

Some writers use these symbols as mere shorthand signs 
for the full names of the elements. This usage is extremely 
objectionable; students who adopt it will not appreciate 
the important quantitative relations which are shown by 
reactions. 

15. Formulae. — Compounds are represented by a formula 
or a group of symbols, showing the composition of the 
molecule of the substance. 

Thus, the formula of sodium chlorid, NaCl, indicates that 
its molecule contains one atom of sodium and one of chlorin, 
and CaS represents a molecule of calcium sulfid which con- 
tains one atom of calcium and one of sulfur. 

If a molecule contains more than one atom of a given 

* The Latin name of sodium is Natrium, that of potassium is 
Kalium. 



10 CHEMISTRY 

element, a subnumber is placed a little below and to the 
right of the symbol, and indicates the number of such atoms. 
Thus CaCL is the formula for calcium chlorid, which con- 
tains one atom of calcium and two of chlorin, and the 
formula for ferric oxid, FeX^, tells us that its molecule 
contains two atoms of iron and three of oxygen. 

If more than one molecule of the substance is to be repre- 
sented, the number is placed before the group of symbols. 
Thus 2 Ft '.,()., represents two molecules of ferric oxid con- 
taining four atoms of iron and six of oxygen. 

In the absence of a coefficient a formula always represents 
a Bingle molecule. 

16. The Law of Constant Proportions. — The law of definite 
proportions which has been called the corner stone of modern 
chemistry i> as follows: — 

" Tin' same compound always contai)is the same elements 
combine I in the same fixed and defi n ite proportions" 

The thousands of analyses which have been made of 
various compounds by chemists in all parts of the world, 
and which are now being made every day, are based upon 
this law. and in no single instance have the results obtained 
caused the truth of the law to be questioned. 

17. Combining Weights. — Another important relation is 
to be learned from a study of the composition of various 
substances. Not only is the proportion by weight in which 
a certain element combines with a certain other element, to 
form a given compound, constant, but it is possible to select 
a number for each element, which shall represent the pro- 
portion by weight in which it unites with different elements. 

The composition of the oxids mentioned thus far is given 
below : — 



SYMBOLS AND LAWS 



11 



Mercury, 200 
Oxygen, 16 

Copper, 63.6 
Oxygen, 16 



Lead, 207 
Oxygen, 16 

Zinc, 65.4 

Oxygen, 16 



Iron, 5\S 
Oxygen, 16 



In each of the above compounds it is observed that there 
are 16 parts by weight of oxygen, and this number or a 
simple multiple of it will express the proportion in which 
oxygen combines with any other element. 

Such numbers have been carefully determined for all ele- 
ments and are called combining weights. 

18. The Law of Multiple Proportions. — The analysis of 
various substances further shows that a given element may 
combine with another in more than one proportion. For 
example, the elements nitrogen and oxygen form several 
compounds having the following composition : — 



Nitrous oxid 
Nitric oxid . 
Nitrous anhydrid 
Nitrogen peroxid 
Nitric anhydrid 



28 parts 
28 parts 
28 parts 
28 parts 
28 parts 



16 parts 
32 parts 
48 parts 
04 parts 
80 parts 



It will be observed that while the quantity of nitrogen is 
the same in all the above compounds the quantity of oxygen 
varies, being twice as great in the second compound as in the 
first, three times as great in the third as in the first, etc. This 
series illustrates the law which applies to all cases in which 
more than one compound is formed from the same elements. 



12 CHEMISTRY 

If two elements form more than one compound, the propor- 
tions by weight in which a given element combines with the 
other in each compound will be expressed either by its com- 
bining number or a simple multiple of its combining number. 

19. The Atomic Theory. — The atomic theory was sug- 
gested by John Dalton, an English schoolmaster, early in 
this century, to account for the laws of definite and multiple 
proportions. It maintains — 

1. That with a few possible exceptions all molecules are 
made up of smaller particles. 

2. That these particles are indivisible (they are therefore 
called atoms). 

.°). That all atoms of a given element are equal in size and 
weight 

4. That atoms of different substances have different 
weigh is. 

5. That the combining weights of the elements are simply 
the relative weights of the atoms, and may therefore be 
called the atomic weights. 

The explanation of the facts of chemistry which this 
theory offers is so satisfactory that it is universally accepted. 

20. Atomic Weights. — As hydrogen enters into combina- 
tion in smaller proportion than any other element, its com- 
bining weight or atomic weight is taken as the unit. When 
we say that the atomic weight of oxygen is 16, we mean 
simply that the atoms of oxygen are sixteen times heavier 
than those of hydrogen. The exact weight of an atom of 
hydrogen has never been determined but it is called a 
microcrith. The atom of oxygen weighs 16 microcriths. 

The following table gives the exact values of the atomic 
weights of the elements referred to in this book. The 
standard is Oxygen = 16. 



SYMBOLS AND LAWS 
Table of Atomic Weights 



13 



Xe.me 


Sym. 


= 16 


Xatue 


Sym . 


= 16 


Aluminum . . . Al 


27.11 


Antimony . . . 


Sb 


120.42 


Argon . 








A 


(?) 


Arsenic 






As 


75.01 


Barium . 








Ba 


137.43 


Bismuth 






Bi 


208.11 


Boron 








B 


10.95 


Bromin . . 






Br 


79.95 


Cadmium 








Cd 


111.95 


Calcium 






Ca 


40.07 


Carbon . 








C 


1201 


Chromium . 






Cr 


52.14 


Chlorin . 








CI 


35.45 


Copper . . 






Cu 


63.60 


Cobalt . 








Co 


58.93 


Gold . , . 






Au 


197.23 


Fluorin . 








F 


19.06 


Iodin . . 






I 


126.85 


Hydrogen 








H 


1.008 


Lead . . 






Pb 


206.92 


Iron . . 








Fe 


56.02 


Magnesium 






Mg 


24.28 


Lithium . 








Li 


7.03 


Mercury . 






Hg 


200.00 


Manganese 








Mn 


54.99 


Nitrogen . 






N 


14.04 


Nickel . 








Ni 


58.69 


Phosphorus 






P 


31.02 


Oxygen . 











16.00 


Potassium . 






K 


39.11 


Platinum 








Pt 


194.89 


Silver . . 






Ag 


107.92 


Silicon . 








Si 


28.40 


Strontium . 






Sr 


87.61 


Sodium . 








Na 


23.05 


Tin . . . 






Sn 


119.0 


Sulfur . 








S 


32.0 


Zinc . . . 




Zn 


65.41 



21. Reaction. — The force which is exerted between atoms 
is called chemical affinity. The affinity of a given atom for 
other atoms varies greatly, often being very strong for cer- 
tain kinds of atoms and feeble for others. If, when any 
substances are mixed, a rearrangement of the atoms would 
produce more stable compounds, i.e. if the force which 
holds the atoms together in the new compounds is stronger 
than that which bound them in their original form, such 
rearrangement will take place. The process of redistribu- 
tion of the atoms in the molecules concerned in the phe- 
nomenon is called chemical action or reaction. 

A reaction is due to chemical affinity and causes a chemical 
change. 



14 CHEMISTRY 

Substances used to bring about desired reactions are 
called reagents. 

The substances which, go into a reaction are called factors, 
and those which come from a reaction products. 

Reactions are ordinarily expressed by equations in which 
the symbols and formulae of the factors are placed on the 
left of the sign of equality, and those of the products on 
the right. The algebraic signs plus and minus are used in 
the ordinary sense in the equations. 

The fact that atoms can neither be created nor destroyed, 
even by chemical means, justifies the use of the sign of 
equality to connect factors and products, and it should never 
be placed until the student has "satisfied the reaction," i.e. 
has determined that there are exactly as many atoms of 
each element in the products as there are in the factors. 

Illustration. In Experiment 11 the following reaction 

occurred : — 

( Ja( Jl s + H 2 S0 4 = CaS0 4 + 2 HC1. 

This should be read as follows : one molecule of calcium 
chlorid plus one molecule of sulfuric acid forms one mole- 
cule of calcium sulfate plus two molecules of hydrochloric 
acid. 

The chemical change occurring in Experiment 19 may be 
expressed as follows : — 

Fe + S = FeS. 

Such equations express very concisely the relations be- 
tween the atoms and molecules in the chemical changes 
which they represent, and every chemical change which is 
clearly understood may be expressed in this way. Equa- 
tions are also useful because of the important quantitative 
relations between masses which are made evident when we 
consider the atomic weights of the elements represented. 



SYMBOLS AND LAWS 15 

In the equation given above, the symbol Fe not only signi- 
fies an atom of iron, but it also stands for 56 parts of iron, 
by weight, and the symbol S stands for 32 parts of sulfur 
by weight. We fchug have a mathematical expression which 
shows the relation between the masses of iron and sulfur 
which take part in the chemical change. 

Fe + S = FeS 
5G 32 88 

The equation can now be read : — 

56 parts of iron unite with 32 parts of sulfur to form 88 
parts of ferrous sulfid. The solution of many chemical 
problems depends upon this use of equations. (See Chap- 
ter IX.) 

22. Analysis, Synthesis, and Metathesis. — All chemical 
changes may be referred to one of four classes : — 

(«) Compound molecules may be separated into their elements, or 
into simpler groups of elements, as, for example, mercury rust is 
separated into mercury and oxygen in Experiment 25, or as potassium 
chlorate KCIO3, is decomposed in Experiment 29, forming potassium 
chlorid KC1, and oxygen. Such changes are analytic, and the process 
which brings them about is known as analysis. 

(b) Compound molecules may be formed by direct union of ele- 
ments, or simpler groups of elements, as when phosphorus combined 
with iodin, in Experiment 10, forming phosphorous di-iodid PI 2 , or 
when carbon monoxid CO, combines with oxygen to form carbon 
dioxid CO2. (See Experiment 96.) Such changes are synthetic, and 
the process is known as synthesis. 

(c) Compound molecules may be formed by a change involving 
both analysis and synthesis, which is known as metathesis, or double 
decomposition. In such processes an exchange of atoms, or groups of 
atoms, takes place between two compound molecules, as when a solu- 
tion of sodium sulfate Na 2 S0 4 , and barium chlorid BaCl 2 , are 
mixed. Each substance is decomposed, and the atoms combine to 
form two new substances, barium sulfate BaSO*, and sodium chlorid 
NaCl. 



16 CHEMISTRY 

(d) In some cases new substances are formed without changing 
either the kinds of atoms, or the number of atoms of each kind, in the 
molecule. For example, when a solution of ammonium cyanate 
XH 4 . O . CN, is heated it is transformed into urea X 2 H 4 CO, a substance 
having entirely different chemical and physical properties. It will be 
observed that these molecules contain the same number of atoms of 
each element ; we have excellent evidence, however, that the first one 
contains cyanogen CX, while the second contains carbon monoxid CO. 

REVIEW QUESTIONS 

1. How many atoms of hydrogen in 6 H2SO4 ? of sulfur ? of 
oxygen ? 

2. How many atoms of each element are represented by the fol- 
lowing formula? : 2 ZnCl 2 , 3 HX0 3 , 5 H 2 0, 14 XH 3 . 

3. How many molecules of each substance are represented by 
above formula? ? 

4. Define chemical affinity, reaction, reagent, factor, product. 

5. Distinguish between atoms and molecules. Does a chemical 
affinity exist between molecules ? Give a reason for your answer. 

6. What is atomic weight ? How is atomic weight related to 
specific gravity ? 

7. What element is selected as the standard of atomic weight ? 
Why is this element selected ? 

8. State the atomic theory. 

9. State'five principles observed in writing chemical symbols and 
formulae. 

10. What is a chemical equation ? What is meant by the combin- 
ing weight of an element ? 

11. State the law of constant proportions. Of multiple proportions. 

12. State five principles to be observed in writing chemical 
equations. 



CHAPTER III 
CHEMISTRY OF THE AIR 

23. The Formation of Rust. 

Experiment XX. — 1. In a small porcelain crucible or a clay pipe 
bowl put a small piece of lead or zinc. Heat with laboratory burner 
and notice the changes that take place. Do not allow the containing 
vessel to become too hot, for liquefie ,1 rust will be absorbed. After the 
lead begins to melt, stir with a thick iron wire. Observe carefully 
what forms on the surface of the metal. Does the lead retain its bright 
mirror-like surface if not stirred ? Continue to heat and stir until the 
substance is changed to a powder. What is its appearance now ? 

2. Let it cool. Is it lead ? What difference is there between the 
action in this case and in nielting ice and cooling the water again ? 
Which is chemical and which is physical action ? Why ? Was the 
change just observed produced by the heat or by the action of the air ? 
In order to answer this question let us repeat the experiment, prevent- 
ing any action of the air by covering the metal with a film of melted 
rosin. 

Experiment XXL — Repeat Experiment 20, adding as much 
powdered rosin as can be lifted on the blade of a penknife. Do 
not stir the metal. Does it rust or does the surface remain bright and 
mirror-like ? Is it changed to powder ? How do you explain the 
difference in result of this experiment and the last ? What do you 
conclude concerning the cause of the change produced in the previous 
experiment ? Is the action due to the high temperature or to the 
action of the air or to both ? Does lead rust more rapidly at high than 
at low temperature ? Rosin is used to prevent rusting of hot metals 
in process of soldering. 

24. Effects of Air on Iron at Ordinary and at High Tempera- 
tures. 

Experiment XXII. — Wind a piece of No. 30 iron wire about a foot 

long around the finger and heat the loops thus formed in the tip of a 

laboratory burner flame for a minute or two. Holding the loop over 

a sheet of paper, straighten the wire. Compare the scale which drops 

• c 17 



18 CHEMISTRY 

off with the rust formed < n iron at ordinary temperatures. Is it the 
same color? Does iron rust more rapidly at high or low tempera- 
tures ? I low do you know ? Has a chemical change occurred ? Pass 
a magnet over a mass of red rust ; of black rust. Are they magnetic ? 
(See paragraph on Oxid.s in Nature, page 35.) 

25. Various Ways of Protecting Iron. 

Several years ago Professor Barff, of London, suggested 
that iron might be protected from the action of the air by 
exposing it to superheated steam at high temperature, thus 
forming a coating of black rust on its surface. The pro- 
cess has been somewhat modified, and is now known as the 
Bower-Barif process. It is quite extensively employed as 
a finish for iron ornaments, and has been used in certain 
cities to protect water pipes. 

Zinc and lead are protected from the action of the air by 
the coating of oxid which forms on their surface. 

Qi ERIE8. — Mention Beveral ways of protecting iron from the action 
of the air. Why do we blacken stoves ? Why are some parts nickel 
plated '. What is galvanized iron ? What is a tin pan made of ? In 
what two ways are water pipes protected ? How are iron bridges pro- 
tected ? Bicycle frames ? 

26. Does the Weight of a Metal change when it rusts ? 
When a chemical change occurs it is due to the addition 

of some element or elements to the substance changed, or 
to the extraction of some element or elements from the sub- 
stance changed. Now, since loss or gain in matter means 
loss or gain in weight, let us determine whether a sub- 
stance was added to or driven off from the iron in the last 
experiment. 

Experiment XXIII. (Performed by the instructor.) — Weigh a 
piece of No. 30 wire, heat as in the last experiment ; when cool weigh 
again. Explain. 

Does heating in contact with air drive something away from the 
iron or cause something to combine with it ? From what source is 
the substance derived ? 



CHEMISTRY OF THE AIR 19 

27. The Material which combines with the Metal to form 
Rust. 

We now desire to know the nature of the substance which 
causes metals to rust ; and as it can be expelled easily from 
the rust which forms on mercury, we shall study that sub- 
stance. As air is an invisible gas, special precautions must 
be taken to prevent its loss or mixture with other substances. 

Experiment XXIV A Method of Collecting Gases. — Fill the 
yellow dish (see description of apparatus, p. 000) one-third full of 
water. Place the flower-pot saucer bottom side up in the water. 
Fi.l aie of the medium sized bottles with water, cover with a glass 
plate, and invert on the flower-pot saucer ; remove the glass plate. 
If your work has been carefully performed your bottle will be full of 
water. (If not, try again.) 

Now put the end of a glass tube at the opening at the s : de of the 
flower-pot saucer and blow gently through it. What do you notice ? 
What is in the bottle after the water is out of it ? Where does it 
come from ? 

This method of collecting gases over water may be used for all 
gases not dissolved by water. 

Students should attempt to devise other methods. Could a rubber 
bag be used ? What advantage has the method used in this experi- 
ment over other methods ? 

Experiment XXV. A Study of Mercury Bust. — 1. Weigh accu- 
rately a small glass tube, closed at one end, containing about a 
gramme of mercury rust. 

2. Holding the tube in a nearly horizontal position with a pair of 
crucible tongs, heat the red powder strongly for some minutes, or until 
a bright mirror-like deposit appears near the open end of the tube. 

3 Weigh the tube again. Is there any evidence that an invisible 
substance has escaped ? After weighing the tube, examine the de- 
posit near the open end. Scrape some of it from the tube with an 
iron wire ; what is it ? What have you learned about the constitu- 
ents of mercury rust ? Is either constituent a solid ? a liquid ? a gas ? 

4. Arrange an ignition tube, as shown in Fig. 1, so that any gas 
generated in the tube may be collected in the bottle. Fill the bottle 
with water. 

5. Put about 15 grammes of mercury rust in the ignition tube and 
apply heat. Describe the gas collected. 



20 



CHEMISTRY 



6. Test with a glowing match stick. Remove the match and put 
it back a few times. Is there any difference between the burning in 
the bottle and out of it ? Is the gas air ? Has the gas which formed 
the rust a marked ability to make things burn ? Has the color of the 
mercury rust changed ? 




7. Remove the ignition tube and pour its contents on a piece of 
paper. How is the color affected ? Compare it with some of the 
mercury rust which has not been heated. What effect has the air 
had upon the hot rust from the tube ? Has the air entirely restored 
the gas driven off by the heat ? Is the gas collected in this experi- 
ment pure air, or a part of the air ? 

The chemical change which occurs in this experiment may be ex- 
pressed as follows : — jjo-0 = H«- + 0. 

The gas which causes metals to rust is called oxygen, its 
compounds are called oxids, and the process of forming oxids 
is known as oxidation. The rust formed on iron at ordinary 
temperatures is called ferric hydroxid. That formed at high 
temperatures is ferrous oxid and ferric oxid, probably in 
chemical combination. It is called magnetic oxid. 

We have observed that oxygen makes things burn vigor- 
ously, and, although it is deemed best to reserve the dis- 
cussion of combustion for a subsequent chapter, the next 
two experiments are given here to show the relation between 
the processes of rusting and burning. 



CHEMISTRY OF THE AIR 21 

28. Effect of excluding the Air from a Flame. 

Experiment XXVI. — Close the holes at the bottom of your labora- 
tory burner. How is the character of the flame affected ? Explain. 
Why do we close the stove dampers at night ? What is the effect of 
removing the ashes and clinker from a stove ? Why ? 

SrcoESTiox. — Wrap a piece of cloth around the lower part of a 
kerosene lamp burner, covering the holes through which the air enters 
the chimney. How does this affect the flame ? What has air to do 
with the combustion of oil ? How does a lamp chimney increase the 
brightness of the lamp flame ? How does a fire extinguisher put out 
a fire ? It is possible that burning, like rusting, is simply a chemical 
union of air, or a part of the air, with the fuel. Let us determine 
whether this is so by the method used in Experiment 23. 

29. Comparison of the Weight of the Products of a Burning 
Candle with the Amount lost by the Candle. 

Experiment XXVII. (Performed by the instructor.) — On one side of 
a delicate balance, apparatus which will absorb the products of combus- 
tion is suspended over a candle, the whole being exactly balanced with 
weights on the other side of the balance. The candle is lighted and 
the gases are drawn into the absorbing apparatus. As the candle 
burns away the side of the balance carrying the apparatus grows 
heavier. The weight of the products is greater than the loss of weight 
of the candle. 

Is the candle converted into heat ? Is heat matter ? Where does 
the matter causing the increase come from ? Does this prove that the 
candle is indestructible ? What is your conclusion concerning the 
nature of combustion ? 

30. Analysis of the Air. — We have learned that oxygen 
is a part of the air, and now desire to learn what proportion 
of the air is oxygen. 

Experiment XXVIII. To determine the per cent of oxygen in the 
air. Cooley's method. Apparatus required. — A small glass funnel. 
A six-inch test tube, with a two-holed rubber stopper to fit same. 
Rubber bands, a measuring glass, 155 cc. of the absorbent liquid. A 
piece of glass tubing two inches long fitted in one of the holes of the 
stopper, a piece of glass rod the same length in the other hole, and a 
piece of thin rubber tubing six inches long, in which a piece of glass 



22 



CHEMISTRY 



C 3=0 , 



Fig. 2. 



rod half an inch long and of such size as to prevent a liquid from 
running through the tube, is placed. 

Manipulation. — 1. Arrange the apparatus as shown in Fig. 2. 

2. In your test bottle dissolve a small teaspoonful 

of pyrogallie acid in 10 cc. of water, quickly add J ^ \ / 
5 cc. of a strong solution of sodium hydroxid, and 
pour into the funnel. This liquid absorbs oxygen 
and carbon dioxid rapidly. 

3. Holding the test bottle under the rubber cork, 
pinch the rubber tube where the glass rod closes it 
until a little of the liquid runs through the tube. 
Carefully remove the drop which is suspended from 
the glass tube with a piece of filter paper. 

\. Now remove the glass rod from the hole in the 
rubber stopper and put the test tube on the stopper; 
allow it to hang there a minute or two to allow the 
heat communicated to the tube and air which it con- 
tains to pass away. Then insert the glass rod in the open hole in the 
stopper. We have now isolated a definite volume of air at the same 
temperature and pressure as the air of the room, and during the ab- 
sorption and the measurements care must be taken to prevent change 
in the volume under analysis, either by the escape of a portion or by 
the introduction of more air from without. 

5. Pinch the rubber tube at the glass rod to allow some of the ab- 
sorbent liquid to run down into the test tube ; a little stream runs in 
at tirst, then drops follow each other more and 
more slowly ; when these have nearly ceased 
allow the apparatus to stand for two or three 
minutes. Then allow more of the absorbent 
liquid to enter the test tube. Repeat the opera- 
tion every two minutes until only a drop or two 
enters the tube when opened. 

6. The gas in the test tube is now compressed 
by the weight of the liquid in the rubber tube ; 
before measurements can be made the pressure 
must be adjusted to that of the air in the room. 
This is accomplished by grasping the test tube 
U_ k by the flange (so as not to warm the gas) , rais- 

ing the tube as shown in Fig. 3, and pinching 
the rubber tube to open a passage between the 
two masses of liquid. Keep this passage open and move the test tube 
up or down until the liquid stands at the same level in the test tube 




Fig. 8. 



CHEMISTRY OF THE AIR 23 

and the funnel ; then close the passage between them. Your results 
will depend to a great extent upon the care with which this adjust- 
ment is made. 

7. Slip a rubber band around the test tube so that its upper edge 
marks the position of the bottom of the stopper. 

8. Remove the test tube from the apparatus and pour the absorbent 
liquid into a measuring glass. This represents the volume of gas 
absorbed ; record the number of cubic centimetres. 

9. Now fill the test tube with water to the top of the rubber band 
and measure this volume. This represents the volume of air analyzed. 

We have thus determined the number of cubic centi- 
metres of oxygen in a certain number of cubic centimetres 
of air, from which we may determine the number of cubic 
centimetres of oxygen in 100 cc. of air; i.e. the percentage 
of oxygen in air. 

31. Other Substances in the Air. — When a gas called 
carbon dioxid is passed through lime water the latter be- 
comes cloudy because a white solid (a precipitate) is formed. 
This is the test for carbon dioxid. 

Experiment XXIX. — 1. Take 20 or 30 cc. of lime water in your 
test bottle. Blow through a glass tube in such a way that the exhaled 
air bubbles through the lime water. Does the lime water become 
cloudy or does it remain clear? What does this experiment prove 
regarding air exhaled from the lungs ? 

2. Eorce air from a bellows through lime water. What inference 
do you draw from this experiment ? 

Carbon dioxid was absorbed with, the oxygen in Experi- 
ment 28, but the amount is so small (about -j^ of one per 
cent) that it may be disregarded. 

Does water vapor exist in the air ? To answer this ques- 
tion, think of the moisture which collects on the outside of 
an ice pitcher in summer. What is dew ? What is frost ? 

The gas which remains in the apparatus after absorbing the 
oxygen and carbon dioxid (Experiments 28 and 30) is nearly 
pure nitrogen. Nitrogen is fully discussed in Chapter VI. 



24 CHEMISTRY 

ARGON 
Symbol A. — Atomic Weight 19.9 

32. Some years ago Lord Rayleigh proved that nitrogen 
obtained by removing the oxygen from the air was invariably 
denser than that obtained from chemical compounds. He 
undertook to determine the cause of this difference, and in 
conjunction with Professor Ramsay found that this greater 
density was due to the presence of an unknown gas, which 
they succeeded in isolating and to which they gave the name 
Argon. Their discovery was announced January 27, 1895. 

Argon is a gas forming T \-$ part of the air; it is also 
found among the occluded* gases in some specimens 
of meteoric iron. As indicated by its name, argon is the 
most inert element ; it has thus far resisted all attempts 
to get it to combine with other elements. Its chief char- 
acteristic, therefore, is its " glorious uselessness." It is 
sparingly soluble in water, boils at — 187° C. and freezes at 
- 189° C. 

Since the discovery of argon several other new elements 
have been found in the air, with properties quite similar to 
those of argon. 

33. Air as a Mixture. — Air is believed to be a mechanical 
mixture of nitrogen and oxygen, and not a chemical com- 
pound, for the following reasons : — 

1. Air contains approximately 79 % of nitrogen and 21 % 
of oxygen. This is not in accordance with the law of 
multiple proportions. 

2. If nitrogen and oxygen be mixed in the above propor- 
tions the mixture possesses all the properties of air, but is 
not accompanied by any phenomena which indicate chemi- 

* Define term. 



CHEMISTRY OF THE AIR 25 

cal action. Whenever chemical union takes place, there is 
some change in the temperature of the substance ; when 
nitrogen and oxygen are mixed as above described there is 
no change in the temperature. 

3. The law of definite proportion states that the compo- 
sition of a given chemical substance is invariable; that of 
air varies slightly. 

4. Air is somewhat soluble in water, but each gas is 
dissolved independently. 

If we shake up air and water in a bottle some of the air 
will be dissolved; if we boil this saturated water the air 
which escapes can be collected and analyzed. This has 
often been done, and it has been found to contain a larger 
proportion of oxygen than the original atmospheric air. 



Thus 


Atmospheric 


Dissolved 


N 


79.04 


66.36 





20.96 


33.64 



This change in the proportion could not occur if the air 
was a compound, for a compound is dissolved as a whole. 
The above numbers exactly agree with the solubilities of 
oxygen and nitrogen separately. 

REVIEW QUESTIONS 

1. Describe the effects of the partial and of the total exclusion of 
air from a flame. 

2. State how the effect of air on iron at high temperatures differs 
from the effect of air on iron at ordinary temperatures. 

3. Describe a chemical method of protecting iron from the action 
of the air. 

4. How does the weight of the products of the combustion com- 
pare with the amount lost by the candle ? Why ? 

5. Does the weight of the scale which flies from the blacksmith's 
hot iron equal the weight lost by the iron ? Why ? 

6. Has air a chemical formula ? 



26 CHEMISTRY 

7. Is the air a mixture or a compound ? Describe an experiment 
to prove the correctness of your answer. Give reasons for your 
answer. 

8. Describe the Bower-Barff process of protecting iron. 

9. Give an account of the discovery of argon. State the prop- 
erties of argon and its occurrence in nature. 

10. Explain the effect of excluding air from a flame. Mention 
some practical appliance whose efficiency depends on the principle 
involved. 

11. What is the scale which accumulates about the blacksmith's 
anvil ? 



CHAPTER IV 
OXYGEN 

Symbol 0. — Atomic Weight 16 

34. Occurrence. — Oxygen is the most abundant of all the 
elements, comprising by weight i of the air, f of the water, 
f of all the animal bodies, and about ^ of the crust of the 
earth. 

The word oxygen means "acid-former," but it is a mis- 
nomer. Chemists supposed that it was present in all acids 
when the name was given. 

35. Preparation. — Oxygen may be easily obtained by 
heating potassium chlorate. 

Caution. — The following precautions must be observed : — 

1. The chemicals must be free from impurities which might cause 
an explosion. If a small quantity of the mixture when heated in a 
dry test tube melts quietly, the mixture may be considered safe. 

2. The ignition tube must be inclined. 

3. It must not be more than one-third full. 

4. The upper part of the mixture in the tube should be heated first. 

5. The heat must be so regulated that an even and not too rapid 
flow of the gas may be secured. It may be necessary to withdraw 
the flame and replace it when the gas slackens. 

Experiment XXX. (Two students will work together. ) — 1 . Arrange 
the apparatus as in Experiment 25. Mix equal weights of manganese 
dioxid and potassium chlorate, and heat about ten grammes of the 
mixture in a test tube. Collect four bottles of the gas evolved over 
water. 

2. Place the bottles on the table, mouth upwards, covering them 
with a glass plate. What is the color of the gas ? Odor ? Taste ? 
Is it soluble in water ? The slight cloud which appears in the, bottle 
27 



28 CHEMISTRY 

at first is clue to a substance which is not oxygen. After a while this 
disappears and oxygen remains. 

3. Drop a piece of charcoal, obtained by charring the end of a 
match stick, in the first bottle. In another lower a deflagrating spoon 
containing a little sulfur. 

4. In the third drop a piece of phosphorus about the size of a 
pea. (Care !) Let them stand quietly and observe what changes, if 
any, take place. Does oxygen at ordinary temperatures act readdy 
on these substances ? 

5. Now thrust a piece of red-hot charcoal (a glowing match stick) 
into the first bottle. Note difference in action. 

6. Remove the deflagrating spoon from the second bottle ; set fire 
to the sulfur. Notice whether it burns with ease or with difficulty. 
Does the sulfur burn more readily in the oxygen than in the air ? 

7. Remove the phosphorus from the third bottle ; place it in the 
deflagrating spoon, ignite, and quickly lower it into the bottle again. 
Describe the action. How does the action of oxygen on these sub- 
stances at high temperatures compare with the action on the same 
substances when cold ? Does either substance burn as vigorously in 
air as in oxygen ? 

Reaction : 2 KC10 3 + Mn0 2 = 2 KC1 + Mn0 2 + 3 2 . 

The Test for Oxygen. — Thrust a glowing splinter of wood into one 
of the bottles. What occurs ? 

Note. — No substance but oxygen can cause a spark to burst into 
flame. How can you determine whether a bottle contains oxygen or 
not? 

36. Physical Properties. — Pure oxygen is colorless, odor- 
less, and tasteless ; it is heavier than air. What are its 
other physical properties ? It is only sparingly soluble, 
water dissolving only 3% of it. Oxygen may be liquefied 
at —118° C. by a pressure of fifty atmospheres. The liquid 
has a pale steel-blue color, and boils at —181° C. under 
ordinary pressure. 

37. Chemical Properties. — Oxygen combines with every 
known substance except fluorin, and is characterized by 
great chemical activity. It is the great supporter of cum- 



OXYGEN 29 

bustiou. If both the oxygen and a combustible substance 
be absolutely dry, it has been shown that the} 7 will not 
combine. No satisfactory explanation of this fact has 
been offered. Oxygen is the only element capable of sup- 
porting respiration. Fish breathe the dissolved oxygen in 
water. 

38. Uses. — Oxygen is necessary to animal respiration, 
to ordinary combustion, fermentation, and decay. It is 
used in the arts to increase the intensity of combustion, 
and is also used in medicine. 

39. Burning in Air. 

Experiment XXXI. — 1. Pour 10 cc. of lime water into a bottle 
containing air, shake the bottle, note the effect on the lime water ; now, 
using a short piece of wire as a handle, lower a burning match into 
the bottle ; when it has gone out cover with the hand and shake the 
bottle ; note the changed appearance of the lime water. A milky- 
appearance proves the presence of carbon dioxid. 

2. Repeat the experiment using a bottle of oxygen. 

When sulfur burns in air a gas having the characteristic odor of 
burning matches is formed. 

3. Determine whether the gas formed when sulfur is burned in 
oxygen is the same that is formed when it burns in air, by burning 
sulfur in a bottle containing air and in one containing oxygen, and 
compare the odors of the gases formed. Discuss the relation between 
combustion in air and in oxygen. 

The difference in activity is due entirely to the fact that 
in air oxygen is diluted with another gas which does not 
support combustion. 

REVIEW QUESTIONS 

1. Describe the preparation of oxygen from potassium chlorate. 
Mention precautions to be observed. 

2. What is the office of manganese dioxid in the above process ? 

3. What are the tests for oxygen ? 



30 CHEMISTRY 

4. Compare the action of oxygen on charcoal at ordinary tem- 
peratures with its action at high temperatures. 

5. Compare the product obtained by burning charcoal in oxygen 
with the product obtained by burning it in air. 

6. Compare the action of oxygen on phosphorus at ordinary tem- 
peratures with its action at high temperatures. 

7. What can you say of the products of combustion in air and in 
oxygen ? 

8. Discuss the occurrence of oxygen in nature. 

9. State the physical properties of oxygen ; the chemical prop- 
erties. 

10. Does oxygen occur uncombined in nature ? 

11. Mention several compounds containing oxygen which occur 
in nature. 

12. Does oxygen display greater energy at high temperatures than 
at low temperatures ? 



CHAPTER V 
COMBUSTION 

40. Ordinary Combustion. — In its broadest sense, the term 
combustion is applied to all cases of chemical action which 
are accompanied by an evolution of heat and light. In the 
majority of cases, however, oxygen is one of the elements 
concerned in combustion, and because of the rarity of the 
exceptions, the term is sometimes defined as the union of a 
substance with oxygen, accompanied by the evolution of 
heat and light ; and the classification of substances as com- 
bustible and incombustible depends upon this definition of 
the term. Thus a combustible substance is one which unites 
with oxygen with evolution of light and heat, and an 
incombustible substance is one which cannot unite with 
oxygen. 

Many substances are products of combustion ; thus water 
is composed of hydrogen and oxygen, and carbon dioxid of 
carbon and oxygen. In these compounds the hydrogen and 
the carbon have already combined with oxygen, and cannot 
directly combine with more. 

41. Blindling Temperature. — A wise provision of nature 
makes it necessary to raise the temperature of substances 
slightly above that which ordinarily obtains, to cause them 
to combine rapidly with oxygen. If this were not true we 
should have no fuels. Substances, differ widely in the tem- 
perature to which they must be raised to cause them to 
combine with oxygen, but for each combustible substance 

31 



32 CHEMISTRY 

there is a definite temperature at which it combines with 
oxygen with sufficient energy to develop heat and light, and 
this is called the kindling temperature. 

If the kindling temperature of a substance is below the 
ordinary temperature, it will take fire when it comes in 
contact with the air, and must, therefore, be kept out of 
contact with the air. Such substances are said to be spon- 
taneously inflammable. Several substances have kindling 
temperatures below a red heat, e.g. the gaseous hydrogen 
phosphid may be ignited with a test tube containing boiling 
water, and the vapor of carbon disulfid may be ignited with 
a glass rod heated to 120°. Most solid fuels require a tem- 
perature slightly above redness, while the diamond must be 
raised to nearly a white heat before combustion begins. In 
starting a fire we take advantage of differences in the 
kindling temperatures of substances. For example, paper 
is easily ignited, but the heat which it develops cannot 
ignite the anthracite ; hence we often put charcoal between 
the paper and the coal, as paper can ignite the charcoal. 
The use of a coating of sulfur or paraffin on matches, to 
.■liable the phosphorus to ignite the wood, is another in- 
stance of the use of a substance having an intermediate 
kindling temperature. 

The temperature produced by the combustion of a sub- 
stance is not necessarily the same as its kindling tempera- 
ture. In all cases of ordinary combustion the temperature 
produced is higher than the kindling temperature of the 
substance; burning particles thus raise adjoining particles 
to the kindling temperature, and the burning continues 
without further application of heat when once started. 

There are, however, numbers of cases in which the tem- 
perature developed by the combustion cannot proceed with- 
out the continuous application of heat. The heat of the 



COMBUSTION 33 

electric spark ignites nitrogen, but the heat developed does 
not kindle the adjacent particles. 

The facility with which a combustible substance may be 
ignited depends upon the quantity of heat, i.e. upon the 
number of heat units required to raise it to its kindling 
temperature. But, as we learn in physics, the temperature 
to which a substance is to be raised is only one of four 
quantities which determine the number of heat units re- 
quired ; the other three being the specific heat of the sub- 
stances, its mass, and the number of heat units lost by 
conduction and radiation. 

The amount of carbon to be kindled in a given stove 
depends upon the specific gravity and the porosity of the 
fuel ; for example, charcoal, gas coke and anthracite coal 
are each of them nearly pure carbon, but they require very 
different amounts of kindling to ignite them. The specific 
gravity of the solid portions of these fuels are as follows : — 

Pine charcoal 40 

Gas coke 86 

Anthracite 1.60 

while the cell space or porosity expressed in cubic centi- 
metres in 100 grammes of the fuel is as follows : — 

Pine charcoal 200.4 

Gas coke 60. 

Anthracite 3.6 

We thus see why charcoal requires comparatively little 
kindling to ignite it, although its kindling temperature is 
the same as the others. 

The amount of heat lost by conduction has an important 
bearing on the amount of kindling required to build a fire. 
If the fuel is a good conductor of heat, it will be diffused 
throughout the mass, and such fuels are more readily 



34 CHEMISTRY 

ignited if they are in small pieces, e.g. shavings are easier 
to ignite than a block of the same kind of wood. 

Experiment XXXII. — Hold the laboratory burner horizontally- 
over a sheet of white paper. Sprinkle some fine iron filings through 
the flame. What occurs ? Pick up a few of the larger pieces on the 
paper and drop them through the flame again. What particles are 
raised to incandescence ? Why ? 

Masses of metal in contact with the fuel often occasion 
considerable loss of heat by conduction. This action is 
nicely illustrated in the following experiment. 

Experiment XXXIII. — Light a candle, bring a piece of wire gauze 
slowly down on the flame until it touches the wick. What occurs? 
Note the conditions above and below the gauze. Hold a lighted match 
above the gauze. What i >ccurs ? Explain. To what extent is the gauze 
heated ? The Davy safety lamp used by miners illustrates this action. 

42. Heat of Combustion. — It must be clearly understood 
that the light produced by combustion is due to the fact 
that the chemical action develops heat more rapidly than it 
can escape, thus raising the body to incandescence. There 
are many cases of oxidation, however, which take place 
slowly, or in which the substance is so situated that the 
heat is conducted away as fast as developed and in which a 
high temperature is not reached. The most important illus- 
tration of this action is the oxidation occurring within our 
bodies, which supplies the heat necessary to our existence. 
Other illustrations are found in the heat developed in com- 
post heaps, in hotbeds, in the decay of wood, in cases of 
" spontaneous combustion," and in the rusting of iron. 

The higher temperature acquired by a substance when it 
burns is readily accounted for by the difference in the rate 
at which it combines with oxygen. When two substances, 
such as carbon and oxygen, combine, their chemical affinity 
causes the atoms to rush toward each other, and the col- 



COMBUSTION 35 

lision which ensues increases their rate of vibration; that 
is to say, it develops heat. The amount of heat developed 
by the collision which forms a single molecule depends upon 
the magnitude of the attractive force and not upon the rate 
at which similar molecules are formed ; it follows, there- 
fore, that the oxidation of a given mass of a substance will 
develop exactly the same amount of heat when it burns that 
would have been developed if it had been oxidized slowly. 

43. Chemical Energy. — All cases of direct chemical com- 
bination are due to attractions between unlike atoms; and 
whether the attraction be great or small, the collison of the 
atoms will develop heat. When molecules consisting of more 
than one atom act upon each other, the force which holds 
the atoms together in the original molecules must be over- 
come before a chemical change can occur; and the amount of 
heat developed in any reaction will accordingly depend upon 
the magnitude of the attractions of the atoms of the factors, 
as compared with, the value of the attractions of atoms of 
the products. If the latter exceed the former, a chemical 
change will occur, accompanied by an evolution of heat. 
Chemical changes which evolve heat are known as ex- 
othermic changes, and those compounds which are formed 
from their elements by such changes are known as ex- 
othermic substances. Such substances are very stable, and 
when they are separated into their original elements the 
same quantity of heat that was evolved when they were 
formed disappears, or more exactly, is transformed into 
chemical potential energy. The formation of a much 
smaller class of substances is accompanied by the dis- 
appearance of heat; these are known as endothermic 
bodies, and when they are decomposed heat is evolved. 
They are usually unstable and often very explosive. All 



36 CHEMISTRY 

endothermic substances possess chemical energy and can do 
work ; that is to say, a substance which can combine with 
other substances without the aid of external energy possesses 
chemical energy. Much of the mechanical energy of the 
world is derived from endothermic substances, e.g. the fuels. 
The decomposition of carbon dioxid in the plant is an 
endothermic reaction in which the energy of the sunlight 
disappears. The carbon thus formed is stored up and may 
be again oxidized. For this reason the energy derived from 
wood and coal is sometimes spoken of as " stored sunlight." 

44. Nomenclature of the Oxids. — The simplest chemical 
compounds are those composed of two elements only; they 
are known as binary compounds. Many binary compounds 
end with the letters id; but this rule cannot be depended 
upon in all instances. 

Binary compounds of oxygen are called oxids ; they are 
very numerous ; e.g. oxygen forms five distinct compounds 
with nitrogen. 

AVI i en there are two oxids of the same element it is quite 
common to distinguish them by adding the suffix ic to the 
name of the element to denote the oxicl having the greater 
amount of oxygen, and the suffix ous to the name of the 
element to denote the oxid having the smaller proportion of 
oxygen. Thus, mercuric oxid has a larger percentage of 
oxygen than mercurous oxid, and nitric oxid a larger per- 
centage than nitrous oxid. 

If there are more than two oxids of the same element, 
prefixes are often used. Thus a peroxid contains a larger 
percentage of oxygen than the oxid to which the suffix ic is 
applied. Nitrogen peroxid, which contains a larger propor- 
tion of oxygen than nitric oxid, illustrates this usage. 

A more scientific and simpler method of naming oxids 



COMBUSTION 37 

has been suggested, and is quite generally used. Accord- 
ing to this plan the first part of the name of the oxid con- 
sists of the name of the element oxidized, and the second 
part of the name indicates the number of atoms of oxygen 
which the oxid contains, by the use of certain prefixes de- 
rived from the Greek. Oxids containing one atom of 
oxygen are called monoxids, e.g. carbon monoxid; those 
containing two atoms clioxids, e.g. carbon dioxid; those con- 
taining three atoms, trioxids, e.g. sulfur trioxid, etc. 

45. Oxids in Nature. — Water, or hydric oxid is the most 
abundant oxid in nature, and sand (silicon dioxid) is next. 

The ores of some of the most important metals are oxids, 
e.g. the red iron ore so common in this country is a com- 
pound of iron and oxygen, the molecule of which contains 
two atoms of iron and three of oxygen ; and black iron ore, 
or lodestone, contains three atoms of iron and four of 
oxygen in its molecule. Many other ores are oxids, e.g. 
those of tin, manganese, etc. 

REVIEW QUESTIONS 

1. How do substances formed by burning in air compare with 
those formed by burning in oxygen ? 

2. Why is not combustion as rapid in air as in oxygen ? 

3. Define combustion. What are combustible substances ? 

4. Define kindling temperature. Which has the highest kindling 
temperature, sulfur, carbon, or phosphorus ? 

5. Mention examples of slow oxidation. How does slow oxida- 
tion differ from combustion ? 

- 6. Compare the amount of heat given off during slow oxidation 
and combustion. 

7. What is meant by chemical energy ? What substances possess 
it ? What substances do not possess it ? 

8. From what source is the mechanical energy of wood derived ? 
Explain. . 

9. What are oxids and how are they named? What do the ter- 
minations " ic " and " ous " indicate ? 



38 CHEMISTRY 

10. What important oxids occur in nature ? Why are they so 
abundant ? 

11. Give evidences that a part of the air combines with the fuel in 
combustion. 

12. Describe an experiment to show the relation of the weight of 
the products of a burning candle to the weight of the portion of the 
candle consumed. 

13. How is combustion related to or distinguished from chemical 
action in general ? 

14. Mention conditions that favor combustion and chemical action 
in general. 

15. Mention a condition favoring some chemical action but not 
combustion. 

16. What is meant by kindling temperature? Explain the theory 
of shaving wood for use in starting a fire of the same kind of wood. 

17. How does the chemical energy of the combustion of hydrogen 
compare with that of the combustion of other elements ? Why ? 

18. Why is a fire of seasoned wood hotte** than a fire of green 
wood '.' 

19. Explain the use of sulfur in making the common friction 
match. 

20. Why is a wood fire easily started with wood shavings? 

21. Upon what does the temperature reached by combustion of a 
given quantity of fuel depend? 

22. Mention live oxids occurring abundantly in nature. 

23. Mention several substances which are acted upon by oxygen at 
ordinary temperatures. 

24. Explain the effect of fine wire gauze when lowered over the 
flame of a lamp. Mention an important practical application of the 
principle involved. 

25. Compare the kindling temperature of hydrogen with that of 
carbon. What bearing has their relative kindling temperature on the 
production of light by illuminating gas ? 

26. Explain the phenomenon of spontaneous combustion. 

27. What would occur if the temperature developed by the com- 
bustion of nitrogen were higher than its kindling temperature ? 



CHAPTER VI 
NITROGEN 

Symbol N. — Atomic Weight 14 

46. Occurrence. — Nitrogen forms 4 of the bulk of the 
air. It is found in combination in a large number of sub- 
stances, e.g. in saltpetre or potassium nitrate, KN0 3 , and 
Chili saltpetre, NaN0 3 . It also occurs abundantly in am- 
monia, nitric acid, flesh, and other animal substances. Its 
compounds give to burned hair and woollens their peculiar 
odor. Many vegetable substances contain nitrogen, as cab- 
bage, mushroom, horse-radish, and it is an essential con- 
stituent of quinine, morphine, prussic acid, and strychnine. 
It forms a part of nearly all explosives, as nitroglycerin, 
gunpowder, etc. 

47. Preparation. — Nitrogen may be prepared by removing 
the oxygen from the air. Any method which burns up the 
oxygen of the air and forms solid or liquid products, yields 
nitrogen which is reasonably pure. If any of the products 
are gaseous they will be mixed with the nitrogen, which 
will therefore be impure. 

First Method. — Introduce a jet of burning hydrogen into 
a bottle of air. After the flame is extinguished there will 
remain in the bottle, nitrogen and the product of the com- 
bustion of hydrogen — H 2 0. 

39 



40 CHEMISTRY 

Second Method. — Phosphorus burns in the air forming 
phosphorus pentoxid, P 2 5 , a flaky white substance which is 
soluble in water. 

Experiment XXXIV. (See Experiment 28). —Support a piece of 
chalk over water in the water pan by means of a wire standard. 
Make a hollow in the chalk, place a piece of dry phosphorus about the 
size of a pea in it. Ignite the phosphorus, quickly cover it with the 
large bottle so that the mouth of the bottle is under water. What 
chemical change takes place ? Notice any change in the volume of 
the air. Explain. Does all the air support combustion ? Take the 
bottle from the water pan (do not allow the water to escape) and 
shake it. What is the result ? Does the white cloud which at first 
filled the bottle remain? Test the gas with a lighted taper. Is it 
combustible ? Is it poisonous ? State its physical properties. 

Third Method. — If air be passed through a tube contain- 
ing heated copper filings, the oxygen combines with the 
copper, forming copper oxide, and nitrogen may be col- 
lected. Nitrogen prepared from the air will contain the 
impurities which exist in the air. Pure nitrogen may be 
prepared as follows : — 

Fourth Method. — Heat ammonium nitrite and collect 
evolved gas over water. 

NH 4 N0 2 = 2 H 2 4- 2 K 

On account of the unstable character of ammonium 
nitrite, it is difficult to keep a supply on hand ; in practice, 
therefore, a mixture of ammonium chlorid and sodium 
nitrite is usually substituted for the ammonium nitrite. 
When this mixture is heated the reaction proceeds accord- 
ing to the following equation : — 

NH 4 C1 + NaN0 2 = XaCl + 2 H 2 + N» 

This method supplies the purest nitrogen. 



NITROGEN 41 

48 Physical Properties. — The last experiment taught us 
certain physical properties of nitrogen; mention them. 
The following additional physical properties not easily 
shown experimentally are worthy of consideration. It is 
sparingly soluble in water, only 1.6 % being dissolved at 
10° C. It may be liquefied at — 193° under pressure of one 
atmosphere. It is slightly lighter than air. 

49. Chemical Properties. — Nitrogen combines directly 
with very few elements, and combination with these ele- 
ments is effected with difficulty. By indirect methods it 
can be made to combine with hydrogen, and with hydrogen 
and oxygen. Its chemical affinities are exceedingly feeble, 
and the compounds which it forms are very unstable. 
Gunpowder, nitroglycerin, and many other explosives are 
nitrogen compounds, and owe their characteristic properties 
to the ease with which they are decomposed. The rapid 
decay of animal and vegetable substances which contain 
nitrogen is a further illustration of the unstable character 
of nitrogen compounds. 

REVIEW QUESTIONS 

1. State the physical properties of nitrogen. 

2. State its chemical properties, activity, combustibility, relation 
to explosives, relation to decay. Is it poisonous ? 

3. What proportion of the air is nitrogen '? How is this shown ? 

4. Describe an experiment for obtaining nitrogen by the use of 
phosphorus. Give the name and formula of the fumes formed, and 
account for their disappearance. 

5. Compare oxygen with nitrogen with respect to (a) chemical 
activity, (&) occurrence, (c) number of compounds, (d) relation to 
combustion and life, and (e) physical properties. 

6. Why is nitrogen an important constituent of most explosives ? 

7. What are nitrogeneous foods ? 



CHAPTER VII 

HYDROGEN 

Symbol H. — Atomic Weight 1 

50. Occurrence. — Hydrogen is never found uncombined 
in nature; its compounds, however, are widely distributed. 
It forms i of the weight of water, and occurs in all 
animal and vegetable matter. It is the only substance 
common to all acids. 

51. Preparation by the Action of an Acid on a Metal. 

Experiment XXXV. — 1. Put a few pieces of granulated zinc in 
the test bottle. Cover them with dilute hydrochloric acid. What 
occurs ? 

2. After a minute or two hold a lighted match over the bottle. 
What occurs ? 




3. Put a few pieces of zinc in your generating bottle. In one hole 
i:i the rubber stopper put a straight glass tube long enough to reach to 
the bottom of the bottle ; in the other fit a bent tube with a delivery 
tube attached. Pour enough dilute sulfuric acid into the bottle to 
cover the zinc. Collect the gas over water. 
42 



HYDROGEN 43 

Caution. — The gas is explosive when mixed with air ; when pure, 
it burns quietly. To determine when all the air which filled the 
bottle at the beginning of the experiment is driven off, collect a small 
bottle of gas ; when full, raise it from the water, mouth downward, 
and apply a match. If the first bottle of gas explodes, repeat until a 
bottle of gas is obtained which burns quietly. 

4. Collect three bottles of pure gas. 

5. Place the first bottle, mouth upward and uncovered, on the 
table. After a few minutes, test it, to see whether or not it contains 
hydrogen. Is hydrogen lighter or heavier than air ? 

6. Pour the hydrogen in the second bottle upward into an inverted 
bottle containing air. Test each bottle with a match. Is there any 
hydrogen in the inverted bottle ? In the other bottle ? 

7. Light a candle with a wire attached for a handle. Hold the 
third bottle mouth downward and thrust the lighted candle well into 
the bottle. What occurs ? What burns ? Does the candle burn ? 
Withdraw the candle slowly. Is it alight ? Why ? Put it back into 
the hydrogen. Does hydrogen support combustion ? Is the mouth 
of the bottle heated ? 

Experiment XXXVI. The Philosopher' 1 s Lamp. (Optional.) — 
Remove the delivery tube and substitute for it a tube drawn out to 
a fine point. If you are sure that the gas is pure, i. e. if you have not 
taken the stopper out of the generating bottle since testing the gas, 
light the gas at the end of the pointed tube. Hold a cold dry bottle 
over the flame. What do you see in the bottle ? Where did it come 
from? (Chemical examination proves it to be pure water). What is 
the product of the combustion of hydrogen ? 

Hydrogen may be prepared by several other processes ; 
for example, by decomposing water by electricity (Experi- 
ment 40) ; by decomposing water by metals at ordinary tem- 
peratures (Experiment 42) ; by passing steam over heated 
metals (Art. 72, etc.). 

The following equations represent chemical changes in 
the preparation of hydrogen by the action of an acid on a 
metal : — 

Zn + 2 HC1 = ZnCl 2 + 2H. 



44 CHEMISTRY 

This is the reaction when hydrochloric acid is used. If 
sulfuric acid is used, the following equation expresses the 
reaction : — 

Zn + H 2 S0 4 = ZnS0 4 + 2 H. 

52. Physical Properties. — Pure hydrogen is odorless, and 
is the lightest known substance ; one litre of it at ordinary 
pressure weighing .08956 gramme. It may be liquefied 
by extreme cold and pressure, but is more difficult to 
liquefy than any other gas. It diffuses more rapidly than 
any other gas. Water dissolves only 2% of hydrogen. 

53. Chemical Properties. — In its chemical affinities hydro- 
; -n closely resembles a metal; it has a strong affinity for 

oxygen, chlorin, and a few other elements, and the com- 
pounds which it forms with carbon indirectly are very 
numerous ; it is, however, very difficult to get it to combine 
directly with carbon. 

54. Comparison of Physical and Chemical Properties of 
Hydrogen and Oxygen. — Hydrogen will burn ; oxygen sup- 
ports combustion. Hydrogen has affinity for few sub- 
stances; oxygen for many. Hydrogen is the lightest known 
substance. Oxygen combines readily with carbon, sulfur, 
phosphorus, and iron. It is difficult to get any of these 
elements to combine with hydrogen. The two elements 
have opposite chemical properties ; yet in their physical 
properties they resemble each other. 

55. Uses. — On account of its great affinity for oxygen, 
hydrogen is extensively used for the purpose of extracting 
oxygen from compounds containing it, i.e. as a reducing 
agent. 

56. Heat and Chemical Energy of the Combustion of Hydro- 
gen. — The chemical affinity of hydrogen for oxygen is 



HYDROGEN 45 

greater than that of any other known substance. The heat 
produced by the combustion of hydrogen is therefore greater 
than that of any other substance. One pound of hydrogen 
in burning gives off 34,400 heat units ; that is, it develops 
enough heat to raise 34,400 pounds of water from 0° C. 
to 1° C. 

The oxy hydrogen blowpipe consists of a tube H (Fig. 5), 
through which hydrogen flows, and at the end of which 
it is ignited. In the centre of this is a small pv tube 




through which a stream of oxygen is forced into the flame. 
The flame produced gives very little light, but its tem- 
perature is between 2000 and 2200° C. ; it is, therefore, used 
in working platinum and other metals fused with difficulty. 
A piece of lime held in the flame is heated to incandes- 
cence, and emits a bright light equivalent to about 120 
standard candles. This device is known as the calcium 
light. 

Experiment XXXVII. — Take notes on the effect of the oxy- 
hydrogen blowpipe flame upon bits of lead, zinc, copper, steel, iron, 
glass, and calcium oxid. 

57. Burning of Oxygen or Air in Hydrogen. — If a jet of 
oxygen or air be introduced into a vessel containing hydro- 
gen, the oxygen or air may be ignited and will burn as 
readily as hydrogen burns in oxygen or in air. 

If a stream of hydrogen be passed through the tube H 
(Fig. 6), and ignited at the bottom of the bottle, a jet of 



46 CHEMISTRY 

air introduced through the hydrogen flame will burn with 
a flame of the same character as that produced when 
hydrogen burns in air. 

58. Product of this Combustion. — The combustion of 
hydrogen must form an oxid of hydrogen. Water is an 
oxid of hydrogen, and analysis of the moisture condensed 
on any cold object held over the hydrogen flame, proves 
that water is the product. 



H 



AIR OR OXYGEN 



59. Formation of this Substance in Ordinary ComDUStion. — 

Nearly all fuels contain hydrogen, and therefore form more 
or less water when they burn ; this can be shown by hold- 
ing a cold object over the flame. Moisture can thus be 
condensed over burning oil, wood, coal, gas, etc. When 
oxygen combines with the waste products of the body 
in the lungs, the hydrogen of the products is turned into 
water; thus the well-known cloud formed by the breath 
in cold weather is this moisture rendered visible by con- 
densation. This may be easily shown by breathing upon 
any cold dry object. 

REVIEW QUESTIONS 

1. Explain the cause of the moisture which appears on a lamp 
chimney when the lamp is lighted. 

2. Why does this film disappear ? 

3. If water is a product of combustion, why does it not extinguish 
the fire ? 



HYDROGEN 47 

4. State the symbol, atomic weight, and occurrence of hydrogen. 

5. Describe the process of preparing hydrogen from zinc and 
hydrochloric acid. Write the reaction. 

6. Discuss the physical and chemical properties of hydrogen. 

7. Describe the oxyhydrogen blowpipe. For what is it used ? 

8. Show how a jet of air may be burned in hydrogen. 

9. What does the moisture which gathers on a cold object held 
over a lighted kerosene lamp indicate as to the composition of the 
kerosene oil ? 

10. Discuss the heat and chemical energy of the combustion of 
hydrogen. 

11. State the color and odor of the gas prepared in Experiment 35. 

12. Is hydrogen explosive ? Under what conditions ? 

13. Does hydrogen combine with oxygen at ordinary temperatures ? 
at high temperatures ? 

14. Describe the hydrogen flame as to (a) color, (6) amount of 
heat, (c) amount of light. 



CHAPTER VIII 

CHEMISTRY OF WATER 

Formula H 2 0. — Molbculab Weight 18 

60. Occurrence of Water in Nature. — Three-fourths of the 
earth's surface is covered with water. It exists in the 
atmosphere, in all vegetable and animal matter, in the soil, 
and even in the rocks. 

/-,>/» riment XXXVIII. — Heat a small piece of alum in a test tube. 
Note evidence that it contains water. Is the loss of water accompa- 
nied by a change in the crystal ? Explain. Repeat the experiment, 
using pieces of gypsum, meat, potato, etc. 

61.. Properties of Water. — At ordinary temperatures pure 
water is a tasteless, odorless, transparent fluid, colorless in 
thin layers, but distinctly blue when viewed in large masses. 
At its greatest density water is 773 times heavier than air. 

Many of the properties of water are used as standards by 
means of which we may express the corresponding proper- 
ties of other substances. The specific gravity of all solids 
and liquids expresses the relation between the weight of the 
substances and the weight of a like volume of water. 

The specific heat of all substances is similarly based upon 
that of water. In the metric sy T stem the unit of weight is 
the weight of a cubic centimetre of water, and the melting 
and boiling points of water are the standard temperatures 
used in the manufacture of thermometers. 

62. Solution. — Water dissolves a greater number of solid, 
liquid, and aeriform substances than any other solvent. 
48 



CHEMISTRY OF WATER 49 

The adhesion between the molecules of the dissolved sub- 
stance and those of water overcomes the cohesion of the 
substance dissolved, and it diffuses through the mass of 
water, forming a transparent solution in which the dissolved 
substance is invisible. Strongly colored substances, how- 
ever, impart their characteristic color to the solution. There 
is a limit to the amount of a substance which a solvent can 
dissolve at a given temperature and pressure, and a solution 
which contains all of a substance which it can dissolve is 
said to be saturated. 

"We can now understand why solution aids chemical 
action. The molecules are no longer held firmly together 
by cohesion; they are free to move, and are thus easily 
brought into the intimate contact necessary to chemical 
action by their chemical affinities. 

If sugar is dissolved in water the solution seems to be 
simply a mechanical mixture ; there is no evidence that a 
chemical change has taken place ; and if the solution is 
evaporated the sugar is recovered unchanged. This, there- 
fore, is a physical solution. 

In Experiment 9, marble was dissolved in hydrochloric 
acid and a chemical change was proven. Such solutions 
are called chemical solutions. 

63. The Effect of Heat on the Solution of Solids and Gases. — 

Physics teaches us that whenever a solid changes to the 
liquid form a certain quantity of heat is rendered latent. 
Hot water can supply this heat more readily than cold, and 
therefore solids are more rapidly dissolved by hot than by 
cold water. Furthermore, a larger quantity of most solids 
is dissolved by hot water than by cold. At high tempera- 
tures the cohesion of the molecules of a solid is less than at 
low temperatures, therefore there is less force to overcome 



50 CHEMISTRY 

in dissolving it. As the water cools, latent heat is with- 
drawn from a certain quantity of the dissolved substance, 
causing it to assume a solid state. 

When a gas is dissolved in a liquid its latent heat must 
be absorbed, but a cold liquid can absorb heat more rapidly 
than a warm one, and therefore a gas is more rapidly dis- 
solved in a cold liquid than in a warm one. As the tem- 
perature of a solution of a gas is raised, a portion of the heat 
is used to change the state of the dissolved gas and a por- 
tion is liberated in gaseous forms. 

64. Water of Crystallization. — Experiment 38 taught us 
that crystalline alum contained water, and that the alum lost 
its geometrical form when the water was driven off. Many 
other crystals are like alum in this respect, and there is evi- 
dence that the water which they contain is held in feeble chem- 
ical combination. The water of crystallization does not make 
the substance moist, as it would if absorbed mechanically, 
and further, a given substance requires a definite amount 
of water for each molecule of the crystal. Certain substances 
form two or more kinds of crystal, requiring different quan- 
tities of water ; and in some crystals the color depends upon 
the amount of water of crystallization. Cobalt chlorid is 
often used as a sympathetic ink because of the change in 
color produced by expelling the water of crystallization. 

In some crystals the water is held so feebly that they lose 
either the whole or a portion of their water of crystallization 
when exposed to the air, and in so doing lose their particular 
geometrical form. This process is known as efflorescence. 

Other crystals absorb water from the air and assume 
other geometrical forms, in some cases absorbing enough 
water to dissolve the crystal. Such crystals are said to be 
deliquescent. 



CHEMISTRY OF WATER 51 

Experiment XXXIX. — Put a crystal' of ferrous sulfate and a 
small piece of calcium cklorid on separate pieces of paper and expose 
them to the air for several days. Which is efflorescent ? Which 
deliquescent ? 

65. Hydroxids. — Strictly speaking hydroxids are com- 
pounds formed by replacing one atom of hydrogen in the 
molecule of water with, an atom of another element or with 
a group of elements. 

Na + H 2 = NaOH + H 
CaO + H 2 = Ca0 2 H 2 . 

According to this definition most acids are hydroxids but 
chemists rarely apply the term to them, whereas, all chem- 
ists agree in calling a compound formed by the union of a 
metal with hydrogen and oxygen an hyclroxid. 

These compounds, which are very important and which are 
discussed more fully in Chapter XII., are sometimes called 
hydrates, but the term is rather objectionable because the 
termination ate is used to distinguish a class of compounds 
to which the hydroxids do not belong. 

66. Electrolysis of Water. 

Experiment XL. (Performed by the instructor. ) — Take notes upon 
this experiment, answering all the following questions and describing 
the apparatus used. An electric current is passed through acidulated 
water from one lead or platinum electrode to another. Gas is evolved 
which is collected in two tubes. Where is the gas liberated ? How 
does the volume over the positive electrode compare with that over 
the negative? What gas is collected over the positive electrode? 
How do you know? Wkat gas is collected over the negative elec- 
trode ? Does this experiment prove that water is composed of two 
elements and no more ? Is the volume of the water decomposed equal 
to the volume of the gases formed ? How do you know ? 

67. Synthesis of Water. 

Experiment XLI. (Performed by the instructor.) — Eudiometer tube 
— a " U " shaped tube of glass about 18 inches long closed at one end, 
having two platinum wires inserted at opposite sides near the closed 



52 CHEMISTRY 

end. Fill the eudiometer tube with mercury and invert in a mercuric 
bath. Introduce a certain amount of oxygen, removing the tube and 
bringing the mercury to a level in both arms. Read the amount of 
oxygen in the tube, filling the arm with mercury again. Introduce 
about twice as much hydrogen in a similar manner, determining the 
exact volume, placing the thumb over the end of the tube that is open, 
and wrapping the tube in a towel, pass an electric shock through the 
wire. An explosion occurs, and water is formed. Some of the gas 
remains in the tube. Testing this residual gas to determine whether 
it is hydrogen or oxygen and subtracting its volume from the quantity 
used, we determine the volume of the two gases which combined. 
This experiment proves that there are only two elements in water. 

68. Formation of Water by passing Hydrogen over Heated 
Oxid. — When mercuric oxid is heated, oxygen is liberated. 
If a stream of hydrogen be passed over a heated oxid, 
the hydrogen and oxygen combine to for.ni water. When 
copper oxid is used, the following reaction takes place : — 

CuO + 2 H = Cu + H 2 0. 
In this experiment if the weight of water formed be de- 
termined and the tube containing copper oxid be weighed 
before and after the heating, it will be found that -| of the 
weight of water came from the copper oxid, thus proving 
that f of the weight of water is oxygen. 

69. Composition of Water by Weight and by Volume. — 
Preceding experiments have shown that water contains 
twice as much hydrogen as oxygen, volume alone consid- 
ered, and that it contains eight times as much oxygen as 
hydrogen, weight alone being considered. These two seem- 
ingly contradictory facts being proven, it follows that the 
single volume of oxygen must be eight times as heavy as 
the two volumes of hydrogen, and that equal volumes being 
considered, the oxygen is sixteen times as heavy as the 
hydrogen. In considering the composition of a compound, 
care must be taken to distinguish between these two 



CHEMISTKY OF WATER 53 

methods of stating the composition, and it must be remem- 
bered that the volumes used are, in all cases, the volume in 
an aeriform state, and not the solid or liquid state. One 
further fact should be stated here : when two volumes of 
hydrogen combine with one volume of oxygen to form 
water they do not form three volumes of steam ; their vol- 
umes is condensed one-third, so that — 

2 vols, of hydrogen -f 1 vol. of oxygen form 2 vols, of steam, 

and this is the way in which the composition of a substance 
by volume should be stated. For further discussion of this 
topic see p. 118. 

The composition of a substance by weight is the same in 
the solid as it is in the liquid and aeriform state. That of 
water may be stated thus : — 

2 parts (by weight) of hydrogen + 16 parts of oxygen form 
18 parts of water. 

70. Decomposition of Water by Metals. 
Experiment XLII. — 1. Fill a medium sized bottle with water and 
invert it in the yellow dish, which should be about half full of water. 

2. Thoroughly dry a small piece of wire gauze in the gas flame, 
roll around a lead pencil, so that it forms a cylinder that is double 
walled in all parts. Fold one end of the cylinder over and pinch the 
bend with a pair of pliers. Drop a piece of sodium the size of a pea 
into it, using the pliers as before. 

3. Lift the bottle slightly, but not enough to allow the water to 
escape, and thrust the wire gauze beneath it. After all action has 
ceased slip a glass plate under the mouth of the bottle and remove the 
bottle from the water pan, placing it right side up on the table. 

4. Test the gas with a burning match. Do you recognize the gas? 
How ? Where did it come from ? Drop a piece of pink litmus paper 
in the bottle. Does it change color ? Does water from the laboratory 
faucet produce a similar change ? What became of the sodium. 

Explain concisely all that has occurred. The reaction is as fol- 
lows : — 

Sodium + Water = Sodium Hydroxid + Hydrogen 

Na + H 2 = NaOH + H 



; "i4 CHEMISTRY 

Note. — Litmus paper is prepared by dipping strips of paper into 
an infusion of litmus. It turns red, when treated with an acid, and 
blue, when treated with an alkali. 

72. Decomposition of Water by passing Steam over Heated 
Metals. — Certain other metals which decompose water 
slowly or not at all at ordinary temperatures, decompose it 
easily at high temperatures. If steam be passed through 
a tube containing bits of iron heated to redness, it will be 
decomposed and hydrogen may be collected over water. 

3 Fe + 4 H 2 = Fe a 4 + 8 H. 

Qi i i:v. — IImw may the weights of hydrogen and oxygen resulting 
from the decomposition be determined in this experiment? 

73. Water Gas. — At high temperatures carbon also de- 
composes water, and this fact is the basis of the process of 
manufacturing water gas. Steam is passed over highly 
heated coal or coke (carbon), which combines with the oxy- 
gen of the water, forming carbon monoxid (CO). The re- 
action is as follows : — 

C + H 2 = CO + 2 H. 
Both carbon monoxid and hydrogen are combustible gases, 
they are both odorless, and burn with feebly luminous flames. 
To overcome these objections it is customary to enrich water 
gas by adding certain gases, obtained by decomposing naph- 
tha, which give the gas a distinct odor, and which greatly 
increase the amount of light produced. The illuminating 
gas used in many of our cities is prepared by this process, 
but the increased value of the ammonia and of the tar 
obtained as by-products in the process of manufacturing 
illuminating gas from bituminous coal has rendered the 
economy of water gas questionable. As water gas is more 
poisonous than ordinary illuminating gas, laws have been 
passed in certain states prohibiting its use. 



CHEMISTRY OF WATER 55 

74. Oxidation and Reduction. — The process of abstract- 
ing oxygen from a body is called reduction. In the manu- 
facture of water gas the hot carbon abstracts the oxygen 
from the water, illustrating its use as a reducing agent. 
Carbon is extensively used in the reduction of ores to a 
metallic state. The union of a substance with oxygen is 
called oxidation, and the reagent which causes the oxida- 
tion is called the oxidizing agent. Nitric acid and potassium 
chlorate are excellent oxidizing agents, as will be noticed 
in several subsequent experiments. In Experiment 42 the 
sodium was oxidized by the water, but water is not among 
the better oxidizing agents. 

75. Natural Waters. — Absolutely pure water is never 
found in nature. The impurities which it contains are of 
two classes : first, the inorganic, or those derived from the 
rocks; and second, the organic, or those derived from the 
decay of animal matter or vegetable substances. Some of 
the impurities are held in solution, while others are sus- 
pended and carried along by moving water. The purest 
water found in nature is rain water, particularly that which 
falls in country districts after it has been raining some 
time. But even rain water contains impurities ; as it falls 
through the air it washes it, removing those suspended 
matters which are always present in it, and dissolving small 
quantities of the gases of the atmosphere. As soon as the 
rain reaches the earth, its great solvent power is exerted 
upon the mineral matter with which it comes in contact, 
and it becomes more impure. An impure water is not neces- 
sarily unfit for household purposes. See Article 78. 

76. Spring Water always contains dissolved mineral mat- 
ter as well as a considerable quantity of carbon dioxid de- 
rived from the decomposition of plants. Waters flowing 



56 CHEMISTRY 

through different strata would naturally contain different 
amounts and different kinds of minerals, and we therefore 
have various kinds of spring waters. Sulfur springs usually 
issue from, rocks containing a decomposing sulfid. A line 
of sulfur springs across western New York marks the out- 
crop of the Hamilton shale which, in certain layers, contains 
a great deal of iron sulfid. The chalybeate springs contain 
some compounds of iron in the same way, and the effervescent 
waters have some gas in solution. The famous springs at 
Saratoga, N. Y., belonging to this class, contain carbon dioxid. 
River water differs from spring water because a part of 
it, at least, has not been filtered through porous rock and 
thus relieved of suspended matter. It is consequently 
turbid, while spring water is usually clear and sparkling. 
Sea-water contains a large amount of mineral matter, the 
average quantity being about 3.5 % of its weight. 

77. Hard Water. — Certain salts which are frequently 
present in natural waters prevent the formation of a lather 
with soap in the ordinary process of washing, and give to 
waters containing them the property known as hardness. 
The chief substances which produce this effect are the com- 
pounds of calcium and magnesium. Soap is decomposed 
by such waters, forming an insoluble, curdy precipitate or 
scum which prevents the cleansing action of the soap until 
all of the hardening salts have been removed. Hardness 
due to the presence of carbonates may be removed by boil- 
ing the water, and is called temporary hardness. Waters 
having this kind of hardness are common in limestone 
regions. The limestone is an impure calcium carbonate 
(CaC0 3 ), and is not soluble in water, but water containing 
carbon dioxid converts the calcium carbonate into an acid 
carbonate (CaH 2 (C0 3 ) 2 ) which is soluble. Permanent hard- 



CHEMISTRY OF WATER 57 

ness, or that which remains after prolonged boiling, is 
usually due to the presence of sulfates. 

78. Potable Waters, — The inorganic impurities found in 
natural waters are very rarely injurious to health. The 
organic impurities, however, are usually accompanied by 
living germs, or bacilli, by means of which such diseases 
as cholera, typhoid fever, diphtheria, etc., are propagated. 
The principal source from which these dangerous organic 
impurities are derived is the drainage of houses and vil- 
lages ; and though waters thus contaminated may become 
pure again through the action of the air and sunlight, it is 
not safe to rely upon this method of purification in the 
water which is to be used for drinking purposes. 

It is of the greatest importance that the water used for 
drinking purposes should be as pure as possible ; to this end 
it should be frequently tested, and if there is the slightest 
suspicion of contamination, it should be thoroughly boiled. 

In general it may be assumed that springs, deep wells, 
and mountain rivers and lakes, are safe sources of water 
supply ; that stored rain water and surface water from 
cultivated land are unreliable ; and that shallow wells and 
rivers, to which sewage gains access, are dangerous sources. 

79. Distillation and other Methods of Purification. — The 

best method of purifying water is by distillation. This 
process removes both organic and inorganic impurities ; and, 
when properly conducted, supplies a perfectly pure water. 
On shipboard the salt water of the ocean is distilled. 

Experiment XLIII. — Using a flask, distil 30 or 60 cc. of water. 
Under "Physical Properties " note color, taste, odor. What action 
has it upon litmus paper ? Does it leave a residue upon evaporation ? 
Answer same questions concerning some natural water. 

Boiling. — Thoroughly boiling water, renders most of the 
organic impurities harmless. Disease germs are destroyed 



58 CHEMISTRY 

and albuminous matter coagulated so that it may be readily 
removed by filtration. 

Filtration does not remove the dissolved inorganic im- 
purities, — these can be removed only by chemical processes 
(see next article) ; but when properly conducted it removes 
both organic and inorganic suspended matter, including 
disease germs. 

Small charcoal or sand filters are more apt to contaminate 
the water passing through them than to purify it ; for after 
a few days' use, the filter becomes so saturated with germs 
that the filtered water contains more of them than it did 
before it was filtered. Unless the filter is so constructed 
that the charcoal or sand may be removed and exposed to 
air and sunlight, it is unsafe to use it. The wire strainers, 
sometimes called filters and made to be attached to faucets, 
may remove some of the suspended matter, but they do not 
remove the germs. The Pasteur filter, in which the water 
passes through a natural stone, is the most efficient small 
filter; the stone must be removed and boiled, or exposed 
to the air for oxidation of the organic matter, occasionally, 
otherwise germs will pass through the filter after a time. 

Chemical Methods. — Attempts are rarely made to purify 
water on a large scale by chemical processes. But such 
processes have been employed for many years by frontiers- 
men and in regions where no drinkable water exists. 

If alum is gradually added to impure water containing 
calcium carbonate, calcium sulfate is formed, carbon dioxid 
is evolved, and aluminum hydroxid is precipitated. The 
aluminum hydroxid entangles the organic matter present 
and settles with it to the bottom, leaving the water clear and 
sparkling. If the water contains an insufficient amount of 
calcium carbonate, it will not be rendered perfectly clear, and 
the deficiency must be supplied by adding sodium carbonate. 



CHE MIS THY OF WATER 59 

Ferric chlorid, iron borings, and potassium permanganate 
have each been used successfully in chemical processes of 
purifying water. 

80. Softening. — (a) Temporary Hardness. 

Experiment XLIV. — 1. Dissolve 1 gramme of soap shavings in 
10 cc. of distilled water in a test tube. Draw out one end of a glass 
tube, about \ inch in diameter, to a point thus : — 



This is to be used as a dropping tube. 

2. Test 10 cc. of hard water, made by passing carbon dioxid 
through lime w r ater until it is clear, in a test tube as follows : Using 
the dropping tube, add a single drop of the soap solution to the hard 
water, shake the tube, repeat the operation as often as may be neces- 
sary to determine the number of drops required to produce frothing. 

3. Test 10 cc. of distilled water in the same way. What do you 
conclude concerning the relative values of hard and pure water for 
washing purposes ? 

4. Boil some of the hard water and test 10 cc. as before. How 
does this compare in value with the unboiled sample ? 

Boiling decomposes the soluble carbonates, expelling 
carbon dioxid and precipitating the insoluble carbonate. 
The "fur" which forms on the bottom and sides of the 
tea-kettle and the scale which forms on the shell and tubes 
of a steam boiler are each due to the repeated removal of 
carbonates from water by boiling. 

Clark' 's process of removing temporary hardness is quite 
extensively used by water-works engineers. The hardness 
is estimated, and the quantity of "milk of lime" required 
to transform the amount of soluble carbonate present into 
insoluble carbonate is then added. The reaction is as 
follows : — 

CaH 2 (C0 3 ) 2 + CaO = H 2 + 2 CaC0 3 . 



50 CHEMISTRY 

(b) Permanent Hardness. 

Experiment XL V. — 1. Test 10 cc. of hard water, made by dissolv- 
ing calcium sulfate in distilled water and filtering, as in Experiment 
41. Note the number of drops of soap solution necessary to produce 
frothing. Boil some of the above-mentioned hard water, and test as 
before. What effect does boiling have on permanent hardness ? 

The sulfates and chlorids of lime and magnesium are 
decomposed by sodium carbonate (common washing soda) 
with the following reaction: — 

CaS0 4 + XaJ '(.), = CaCO., + Xa,S0 4 . 
The sodium sulfate produced has no effect on the soap, but 
water containing it should not be used for drinking pur- 
poses. Hence the use of washing soda for softening water. 

81. Natural Methods of Purification, (a) Action of the 
Air. — Disease germs die when exposed to the sunlight, and 
organic impurities are oxidized when exposed to the air. 
Therefore, impure water running through shallow streams, or 
over precipices in a thin sheet, tends to become pure again. 

(b) Filtration through beds of sand or porous rock 
removes suspended matter. It is in this way that most 
spring waters are rendered clear and transparent; it should 
be remarked, however, that clearness is not unfailing evi- 
dence that water is healthful. 

(c) Sedimentation. — In ponds and lakes water is often 
rendered clear, the suspended matter settling to the bottom 
under the influence of the force of gravity. 

(d) Natural Distillation. — This is the most important 
method of purifying the water in nature. The sun is the 
source of heat, and since evaporation takes place at a 
temperature far below the boiling point, it occurs every- 
where and at all times, even in winter, and the amount 
evaporated every hour is enormous. The vapor rises to 
the upper atmosphere, and encountering cold currents is 



CHEMISTRY OF WATER (31 

condensed in microscopic particles which float in the air. 
These increase in size as they move about, and finally fall 
as rain. The purest water found in nature is rain water, 
particularly after it has rained some time. 

Mountain streams which flow over rocky beds, notably 
those which flow over beds of sandstone, have exceptionally 
pure waters. Water which flows over limestone dissolves 
some of the stone and becomes hard. 

Note. — It is suggested that students take advantage of the oppor- 
tunity offered by Experiments 46-48 to examine the water which they 
habitually use for drinking purposes. 

82. Tests for Organic Impurities. 

Experiment XLVI. — Fill a tall glass jar with water to be tested. 
Add a few drops of sulfuric acid, then add a weak solution of potas- 
sium permanganate, drop by drop, until the water assumes a violet 
tint. If organic matter be present, the color gradually grows lighter. 
If the color remains unchanged for an hour, the water may be con- 
sidered safe for drinking purposes. 

Xessler's Test. — Nessler's reagent is prepared by mixing potassium 
iodid and mercuric chlond, and adding caustic soda. It furnishes 
a very delicate test for free ammonia, which is evidence of decom- 
posing organic matter. 

A drop or two of the reagent added to water containing ammonia 
gives it a brown color ; the greater the amount of ammonia, the darker 
the shade of brown. In practice it is customary to concentrate the 
ammonia in 500 cc. of the water to be tested by distilling it with a 
small quantity of sodium carbonate, and testing the first 50 cc. of the 
distillate with ISTessler's reagent. 

Experiment XL VII. Test for Chlorids. 

Note. — The members of the class should secure samples of drink- 
ing water from as many sources as possible, including water from a 
shallow well, a deep well, a pond, and a stream. Distilled water and 
a solution of common salt will be required ; these are for comparison 
only, as distilled water contains no chlorin, and salt water a great 
deal. (Common salt is composed of sodium and chlorin. ) 

1. Add a few drops of nitric acid, free from chlorin, to 25 cc. of 
the distilled water, then add a few drops of silver nitrate solution. 

2. Treat 25 cc. of the salt water in the same way. Compare this 



62 CHEMISTRY 

with the distilled water. How can you distinguish water containing 
chlorin from pure water? 

:!. Concentrate 50 cc. of the drinking water to be tested to 25 cc. 
by boiling, and repeat the test. 

Docs it contain chlorin? Compare your result with that of other 
members of the class, and answer the following questions: — 

Does the water from the shallow well become milky? Does that 
from the deep well ? Which samples contain chlorin ? 

Sewage always contains chlorids, and hence if a drinking 

water is found to contain a chlorid, it is to be suspected, 
and the chlorin must he proved to come from some other 
source, or the water should be avoided. If a well is sunk 
near the sea, or near a deposit of rock salt, its water may 
be perfectly wholesome although containing a relatively 
large amoum of chlorids. 

83. Tests for Inorganic Impurities. — The presence of 
inorganic impurities is usually made known without a cheni- 
ieal tesl : the presence of hydrogen sulfid is detected by 
the odor, free carbon dioxid by effervescence, iron by the 
taste, and dissolved salts by the action of soap. 

The presence of solids in solution may be determined by 
evaporating a few drops of the water to be tested on clean 
platinum foil. If solids be present, a residue will remain 
on the foil. 

The fact that one usually desires to know concerning the 
inorganic impurities of w r ater supplied for household use 
is its degree of hardness. 

Experiment XL VIII. To determ im the Degree of Hardness. — Pour 
70 cc. of the water under examination into a flask. Add 1 cc. of 
•'Clark's soap solution," insert a stopper, and shake thoroughly. 
Set it aside for two or three minutes ; if a lather does not remain on 
the surface of the water at this time, add a second cubic centimetre 
of the soap solution. Repeat this process until a permanent lather is 
obtained. The number of cubic centimetres of soap solution used 
is equal to the number of degrees of hardness, and is one greater than 
the number of grains of calcium carbonate per imperial gallon. 



CHEMISTRY OF WATER (33 

84. Hydrogen Dioxid, IL0 2 . — This is a colorless bitter 
liquid somewhat heavier than water. It is very unstable, 
the molecule breaking up into water and oxygen. Because 
of this property it is manufactured on a large scale for use 
as a bleaching agent. It is also used as a disinfectant, and 
to some extent in laundries ; it bleaches the skin and hair, 
and rapidly oxidizes metals. 

REVIEW QUESTIONS 

1. Describe the compounds which hydrogen forms with oxygen. 

2. Describe two methods of decomposing water, and state how 
the weight of each gas may be determined in each case'. 

3. Mention three ways in which water may be decomposed, and 
two ways in which water may be formed. 

4. Describe the electrolysis of water, and state how yon would 
determine which gas is oxygen. 

5. Describe an experiment illustrating the preparation of water 
by passing hydrogen over heated oxids. 

6. What is an analytic experiment ? — a synthetic experiment ? 

7. Describe experiments which prove (a) that water is composed 
of oxygen and hydrogen, (b) that it contains no other element, 
(c) that it contains twice as much hydrogen as oxygen by volume. 

8. Describe the manufacture of water gas. Give equation ; com- 
pare it wuth other illuminating gas. 

9. Distinguish between an oxidizing agent and a reducing agent. 
Give examples. 

10. Describe an experiment illustrating («) oxidation (6) reduction. 

11. What is meant by the term "hard water"? Explain how 
hard water may be made soft. 

12. Distinguish between permanent and temporary hardness. 

13. Give at least two tests by which impurities may be detected in 
water which appears pure to the eye. 

14. How may water containing organic impurities be so purified as 
to render it safe to use ? 

15. Describe three ways in which water is purified by nature. 

16. Describe the Pasteur water filter, and compare it with some 
other filter as to efficiency and as to necessary care. 

17. Account for the formation of caves in limestone regions. 



CHAPTER IX 
PROBLEMS 

85. Composition by Weight. — For the present purpose 
we will assume that the formulas of many substances, not 
yet examined by the pupil, are known. The formula of 
each compound shows the kind of atoms and the number 
of each kind in a molecule of the substance. Considering 
the atomic weight of each element, we may determine from 
the formula the number of parts, by weight, of each ele- 
ment in the compound. Thus, in the case of water, H 2 0, 
the formula tells us that each molecule contains two atoms 
of hydrogen and one of oxygen ; the two atoms of hydrogen 
weigh two microcriths, and the atom of oxygen weighs 
L6 microcriths; the molecule of water, therefore, weighs 
18 microcriths, and T 2 g of the weight of the molecule is 
hydrogen, and || oxygen; further, since all molecules are 
exactly alike, -fa of the weight of any number of molecules 
of water must be hydrogen, and }§ oxygen. 

The molecular weight of any substance is the sum of the 
weights of all the atoms in the molecule. .Therefore, to 
find the molecular weight of a given molecule, add the 
number of microcriths of each element composing the 
molecule. 

Case I 

86. To compute the Weight of an Element in a Given 
Weight of a Compound. — To solve these problems it is 
merely necessary to know how many microcriths of that 
64 



PROBLEMS 65 

element there are in the total number of microcriths in the 
molecule ; since the ratio of the weights in the total amount 
of the substance is the same as in the molecule. 

Example. — How much oxygen in 100 lbs. of sulfuric acid ? 

From the formula of sulfuric acid, H 2 S0 4 , we can (knowing the 
atomic weights of each of the elements) calculate that the molecule 
weighs 98 microcriths, of which (34 microcriths are oxygen. The 
oxygen is therefore f| of the whole weight, and ff x 100 = 65.31 lbs., 
which is the required answer. 

Problems may be solved by proportion as follows : — 

Molecular weight : microcriths of element 
= weight compound : weight of element. 

In above example we have : — 

98 : 64 = 100 : x, 

x = 65.3 lbs. Ans. 

Rule. — Determine the number of microcriths in the mole- 
cule of the compound. Multiply the given weight of the com- 
pound by the number of microcriths of the element divided by 
the total number of microcriths in the molecule, and the result 
is the quantity of the element in the given weight of the com- 
pound. 

Case II 

87. To compute the Weight of a Substance Required to 
Combine with, or Obtainable from, a Given Weight of Another 
Substance. — If substances always combined molecule for 
molecule, the weight of each substance required would be 
proportional to the molecular weights ; but as this is the 
exception rather than the rule, it is necessary to consider 
the reaction to determine the number of molecules of each 
substance required. In the preparation of hydrogen, the 
following reaction occurs : — 

Zn + 2 HC1 = ZnCl 2 4-2H. 



i; ( ; CHEMISTRY 

It will be observed that two molecules of hydrochloric 
acid are required to form one molecule of zinc chlorid. 
Now one molecule of hydrochloric acid weighs 36.5 micro- 
criths, and two molecules 73 microcriths. One molecule of 
zinc chlorid weighs 65 + 35.5 + 35.5 = 136 microcriths. 
Therefore 73 microcriths of hydrochloric acid will t form 
136 microcriths of zinc chlorid, and larger weights of the 
two substances will be in the same proportion; hence the 
following : — 

Rule. — Write the reaction. Determine the number of 
microcriths of each substance involved in the problem. TJie 
ratio of the microcriths wiU equal the ratio of the actual 
weights. 

Example. — 1. What weight of hydrochloric acid can be made 
from 100 lbs. of sodium chlorid? 

Reaction: 2 NaCl + H 2 SO = NaaSOi + 2 HC1. 

2 molecules of NaCl weigh 117 microcriths. 
2 molecules of HC1 weigh 73 microcriths. 
117 :73 = 100 :x, 
117 x = 7300, 

X = 62.3 lbs. Ans. 

2. How much oxygen required to burn 10 grammes of carbon ? 
C + 2 O = C0 2 , 
12 : 32 = 10 : X, 
\2x = 320, 

x = 26| grammes. Ans. 

Case III 

88. To find the Weight of a Litre of Any Gas. — In physics 
we learn that the specific gravity of a substance is the ratio 
of the weight of the substance to that of a like volume of 
some standard. Two standards are in common use, the air 
and hydrogen, and to avoid confusion it is customary to 
apply the term vapor density to the weight of a gas as com- 

L,tfG. 



PROBLEMS 67 

pared with hydrogen. The vapor density of a gas, there- 
fore, simply expresses the fact that any volume of the gas 
is just so many times as heavy as a like volume of hydrogen 
under like conditions of temperature and pressure. The 
letters N.T.P. (normal temperature and pressure) are often 
used in connection with problems of this kind. One litre 
of hydrogen, N.T.P., weighs .0896 gramme. This quantity 
is called a crith. 

The vapor density of each element thus far considered is 
equal to its atomic weight. The vapor density of any gas 
is equal to one-half its molecular weight. 

Example. — To find the weight of 10 litres of oxygen. 

From the table on p. 14 we find that the atomic weight of oxygen 
is 16. As it is an element, this number is also its vapor density ; 
hence 10 litres will weigh .0896 gramme x 10 x 16 = 14.33 grammes. 

Ans. 

Rule. — Multiply the iveight of one litre of hydrogen by the 
vapor density of the gas. 

Case IV 

89. To compute the Percentage Composition of Any Sub- 
stance. — Express as a decimal the ratio of the weight of 
each element to the weight of the compound. These deci- 
mals give the amount of each element in a unit weight of 
the compound, and in 100 weights there would be 100 times 
as much of each element. Move the decimal point, in each 
case, two places to the right, and the percentage composi- 
tion is therefore obtained. 

Note. — A litre of air weighs 14.44 criths. 

REVIEW QUESTIONS 

1. How many kilos of sulfur are contained in 49 kilos of sulfuric 
acid, H 2 S0 4 ? 

2. How many grammes of carbon dioxid, C0 2 , should be obtained 
by burning a diamond weighing \ gramme ? 



Q8 CHEMISTRY 

3. Find the weight of oxygen in 100 grammes of ferric oxid, Fe 2 3 . 

4. How much hydrochloric acid, HC1, can be made from 100 lbs. of 
salt, NaCl ? 

5. If 100 grammes of oxygen are to be converted into water, how 
many grammes of hydrogen must be used ? 

6. How many kilos of hydrochloric acid may be obtained from 20 
kilos of salt ? 

7. How much potassium chlorate, KC10 3 , would be required to 
evolve 2 lbs. of oxygen ? 

8. In 126 lbs. of nitric acid, HN0 3 , how much oxygen ? 

9. Find the proportion of each element in chloric acid, HC10 3 , by 
volume. By weight. 

10. How many pounds of hydrochloric acid, HC1, can be made 
from the chlorin found in 234 lbs. of sodium chlorid, NaCl ? 

11. How many kilos of sodium chlorid would be required to pre- 
pare 73 kilos of hydrochloric acid ? 

Reaction : 2 NaCl + H 2 S( ) 4 = Na 2 S0 4 + 2 HC1. 

12. How many pounds of sulfuric acid will be required to convert 
1010 lbs. of potassium nitrate, KN0 3 , into nitric acid, HN0 3 ? 

Reaction : 2 KN0 3 + H 2 S0 4 = K 2 S0 4 + 2 HN0 3 . 

13. Find the number of grammes of sulfuric acid required to fur- 
nish 100 grammes of hydrogen, by its action on iron. 

14. If 1 litre of oxygen weighs \\ grammes, how many litres of 
oxygen will result from the decomposition of 1 kilogramme of water ? 

15. If a liter of hydrogen weighs .089578 gramme and a litre of 
oxygen weighs 1.429 grammes, what is their proportion by weight in 
water ? 

16. If 806.4 grammes of water are decomposed, what is the volume 
of each of the resulting elements ? 

17. Find the number of grammes of potassium chlorate required to 
furnish 100 litres of oxygen. 

18. Find the number of litres of hydrogen that may be set free by 
the action of sulfuric acid on 10 grammes of zinc. 

19. Find the number of litres of oxygen obtained from 100 grammes 
of potassium chlorate heated with manganese dioxid. 

20. Compute the percentage composition of potassium chlorate. 
How many grammes of oxygen would be given off by heating 500 
grammes of the chlorate, and what would be its volume ? 

21. Find the percentage composition of water by weight and by 
volume. 



PROBLEMS (39 

22. A litre of water weighs 1000 grammes ; how many grammes of 
hydrogen does it contain ? how many grammes of oxygen ? how many 
litres of each ? 

23. How many litres of the element in the first column may be 
obtained from 10 grammes of the substance in the second column '? 

H H3SO4 

H0SO4 

HNO3 

H HN0 3 

N HNO3 

H HC1 

CI HC1 

24. How many grammes of the substance in the first column will 
be required to prepare 10 litres of the element in the second column ? 

H 2 S0 4 H 

KCIO3 CI 

KCIO3 

Zn H 

H 2 ( ) O 



CHAPTER X 

COMPOUNDS OF NITROGEN AND HYDROGEN 

Nitrogen combines with hydrogen to form three com- 
pounds; namely: — 

Ammonia, XHo 

Hydrazine, NgH 4 

Hydrazoic acid, N 3 H 

AMMONIA 
Formula NHg — Moleci lai: Weight 17 

90. Occurrence. — Ammonia occurs in small quantities 
in the air, being formed under certain conditions by the 
decay of animal and vegetable substances. 

The chief source of ammonia is the ammoniacal liquor 
of the gas works, which is the water through which the 
gas has been passed to remove the ammonia formed by 
the decomposition of the coal. 

91. Preparation of Ammonia, NH 3 . 
Experiment XLIX. (For two students.) 

Note. — Experiment 50 may be performed with this one if the 
students will arrange the three bottles with tubes required before 
beginning this experiment. 

1. Mix 4 grammes of ammonium chlorid and 8 grammes of calcium 
hydrate on a piece of glass, adding a few drops of water. Place the 
mixture in a flask with a suitable delivery tube, add sufficient water 
to cover the mass, and apply heat. 

2. Collect three bottles of the gas by upward displacement. Set 
them aside, mouth downward. 

70 



COMPOUNDS OF NITROGEN AND HYDROGEN 71 

3. Hold the Bunsen burner flame in a stream of the gas. Can you 
ignite it ? Does the gas burn while the burner is held in the stream ? 
Is the gas combustible ? Can its flame be said to have a color ? 

4. Connect the flask to the series of bottles described in Experi- 
ment 50, and proceed with that experiment. 

5. When leisure permits, test the three bottles of the gas in such 
manner as to enable you to answer the following cpaestions : — 

Is it a supporter of combustion ? Is it heavier or lighter than air ? 
Is it soluble in water ? Test with a piece of moist pink litmus paper. 
Note the odor. Caution. 

When ammonia is prepared as above, the action is as 
follows : — 

CaHA + 2 XH 2 C1 = 2 NH 3 + CaCL, + 2 H 2 0. 

Ammonia is also formed when nitrogenous organic matter 
is heated out of contact with the air, as in the process of 
making illuminating gas by heating coal, or in the process 
of making animal charcoal. 





Ammonia gas is often prepared by heating the stronger 
ammonia water of commerce in a flask and collecting the 
gas as in Experiment 49. 

Ammonia water was formerly called spirits of hartshorn, 
because it w r as prepared by distilling the horns of the 
hart. 



72 CHEMISTRY 

92. Preparation of Ammonium Hydroxid, NH 4 OH (Ammonia 
Water). 

Experiment L. (For two students.) — 1. Connect a series of 
three medium sized bottles with the delivery tube of the flask used 
in Experiment 4 ( .t as shown in Fig. 8. Have 20 cc. of water in the 
first bottle, and about 50 cc. in each of the others. The first and 
second bottles are fitted with rubber stoppers. Neither tube of the 
first bottle should dip below the water. The tubes by which the gas 
enters the remaining bottles should dip beneath the water. The gas 
is sometimes dissolved faster than it is supplied in a given bottle, and a 
vacuum is formed, causing the water to run into it from the next bottle- 
Should this occur, raise the stopper of the bottle toward which the 
water flows. It three-holed stoppers are at hand, it is best to use 
them, and to insert safety tubes in the first and second bottles, thus 
preventing this action. 

2. Hold a glass rod moistened with hydrochloric"acid over a bottle 
of ammonia water. What occurs ? This is the test for ammonia. 
Tot the ammonia water with litmus paper. 

93. Manufacturing Processes. — In the arts, ammonia is 
prepared either by adding slaked lime to the ammoiiiacal 
liquor of the gas works, or by boiling the liquor. In 




the latter process the vapor is often passed into sulfuric 
acid, forming crude ammonium sulphate, which is used 
as a fertilizer. 

At the Eochester Ammonia Works the liquor is forced 



COMPOUNDS OF NITROGEN AND HYDROGEN 73 

into the generator G, Fig. 9, by a steam siphon S, and 
slaked lime is introduced through the pipe L. The mix- 
ture is agitated and heated, and the gas escapes through 
the delivery pipe D to the iron tanks C, where it is dis- 
solved in water as in Experiment 50. Between G and 
C, the gas passes through several cylinders containing 
petroleum and various other solvents which absorb im- 
purities, but allow the ammonia to pass through them. In 
this way, the " stronger ammonia water " of the drug stores 
is made. 

94. Liquid Ammonia. — There is a large demand for 
anhydrous ammonia for use in ice machines. It is usually 
prepared as follows : Ammonia water is heated in an iron 
cylinder A, Fig. 10. The 

gas, NH 3 , is driven over 
into the condenser C. As 
it accumulates, the press- 
ure increases until it 
reaches 130 pounds per 
square inch, when the gas 
is liquefied. 

95. Physical Properties. — Ammonia gas can be easily con- 
densed to the liquid form by cold and pressure. When the 
pressure is removed it passes back to the form of gas and 
absorbs heat in so doing. 

One volume of water dissolves 600 volumes of ammonia 
gas at ordinary temperatures. It condenses at ordinary 
temperatures at 6.9 atmospheres; at ordinary pressures it 
boils at — 33.7° C. and solidifies at — 75°. 

96. Chemical Properties. — A jet of ammonia to which a 
flame is applied continues to burn in oxygen after the 
flame is withdrawn. In air, however, the heat of com- 



A 




C 


Fig. 10. 




gjJ 





74 CHEMISTRY 

bustion is not sufficient to raise adjoining particles to the 
kindling temperature. 

In gaseous form as well as in the solution ammonia turns 
red litmus paper blue and neutralizes acids just as the alka- 
lies, sodium, and potassium hydroxids do. 

Ammonia combines directly with acids and many other 
substances to form a series of compounds which resemble 
each other in general properties, and each of which contains 
the group of atoms ,XH 4 . In its chemical action, therefore, 
this group resembles a metallic element, and in order to dis- 
tinguish it from ammonia, NH 3 , the termination urn, which 
is applied to nearly all metallic elements, is substituted for 
the final a of ammonia. The similarity in composition of 
some of the compounds of ammonium, zinc, and sodium is 
shown in the following table : — 





NU 4 


Zn 


Na 


NH4CI 

(NH 4 ) 2 S0 4 
NH4NO3 


ZnCl 2 
Z11SO4 

Zn(N0 3 ) 2 


NaCl 

Na 2 S0 4 
NaNOs 





Compounds which play the part of an element are called 
radicals. 

There is more or less evidence that a definite chemical 
compound is formed when ammonia is dissolved in water 
and that its formula is NH 4 OH. This compound has never 
been isolated, however, and certain chemists doubt its exist- 
ence. The name ammonium hydroxid implies that the 
solution is chemical rather than physical. 

97. Uses. — Ammonia is extensively used in the manu- 
facture of artificial ice, aniline colors, indigo, washing soda, 



COMPOUNDS OF NITROGEN AND HYDROGEN 



75 



etc. It is also used in the laboratory as a reagent and in 
the household as a detergent. 

The following device illustrates the manner in which 
liquid ammonia is employed in making artificial ice : — 

Anhydrous ammonia flows from the tank LA into the 
chamber which surrounds a vessel of water i"; it evapo- 
rates rapidly, and so much heat is rendered latent that the 
water is frozen. The gas is pumped back into the tank 
and again liquified by pressure, and the operation is re- 
peated. 

In the Carre Ice Machine, a pound of coal is consumed 
for each pound of ice made. 



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REVIEW QUESTIONS. 



1. Describe the preparation of the ammonia of commerce. 

2. Describe the preparation of ammonia in the laboratory. 

3. Discuss the combination of ammonia with water. 

4. Distinguish between ammonia and ammonium ; between liquid 
ammonia and ammonia water. 

5. In what respect does ammonium resemble a metal ? 

6. What is hartshorn ? 

7. State the color, odor, solubility, and weight of ammonia. 



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